LABOMTORY  PRACTICE 
N  CHEMISTRY 


UC-NRLF 


SB    3Db    552 


GIFT  OF 
Publisher 


EDUCATION  DEPT. 


From  the  collection  of  the 


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Prelinger 

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JJibrary 

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San  Francisco,  California 
2006 


Funnel 


Mortar  and  pestle 

Pipe  stem  triangle 


Forceps 


Deflagrat  ing 
Funnel  tube 


Blowpipe 


Florence  flask 


Evaporating  dish 
Beaker 


Porcelain  crucible 
and  cover 


Wa 

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tch  or  cover 
ass 


30cc.  graduated 
test  tube 


File 


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Bunsen  burner 


Iron  ring  stand 


Test  tube  rack  with.  ^ 

hard  glass  test  tubelA  \and  ordinary 
test  tube[B\  n 


Some  apparatus  included  in  the  student's  outfit 


LABORATORY   PRACTICE 
IN   CHEMISTRY 

EXERCISES  TO  ACCOMPANY 

"CHEMISTRY  AND  ITS  USES" 

BY 

WILLIAM  McPHERSON 
it 

AND 

WILLIAM  EDWARDS  HENDERSON 


GINN  AND  COMPANY 

BOSTON    •    NEW    YORK    •    CHICAGO    •    LONDON 
ATLANTA    •    DALLAS    •    COLUMBUS    •    SAN    FRANCISCO 


D45 

Mi, 


COPYRIGHT,  1922,  BY 
WILLIAM  MCPHERSON  AND  WILLIAM  EDWARDS  HENDERSON 


ALL   RIGHTS   RESERVED 
S23.9 

Qfft     Published 
EDUCATION  DEPT. 


gftc   fltftengum 

GINN  AND  COMPANY  •  PRO^ 
PRIETORS  •  BOSTON  •  U.S.A. 


auc 


PREFACE 

In  selecting  the  experiments  to  accompany  the  text  "Chem- 
istry and  its  Uses"  the  authors  have  made  every  effort  to  limit 
their  choice  to  such  as  have  been  subjected  to  extended  and 
successful  trial  in  a  large  number  of  secondary  schools  in  which 
ordinary  conditions  prevail.  The  authors  have  relied  not  only 
upon  their  own  experience  with  large  numbers  of  students  in 
elementary  chemistry  but  also  upon  the  cooperation  and  crit- 
icism of  a  number  of  teachers  in  secondary  schools  throughout 
the  country.  It  is  a  pleasure  to  acknowledge  in  this  place  the 
assistance  so  willingly  rendered  us. 

It  is  clearly  impossible,  however,  to  make  a  selection  that  will 
meet  with  the  approval  of  all  teachers.  Indeed,  an  author  can 
hardly  hope  to  secure  the  entire  approval  of  any  one  teacher. 
Each  one  works  under  certain  limitations  of  time,  equipment, 
and  local  interests,  and  each  has  his  individual  training  and 
viewpoint.  An  experiment  regarded  very  highly  by  one  is 
often  considered  rather  fruitless  by  another.  One  desires  to 
stress  the  applications  of  chemistry  and  another  prefers  to  ad- 
here to  more  fundamental  principles.  And  all  this  is  exactly 
as  it  should  be  with  enthusiastic  teachers. 

In  order  that  each  teacher  may  have  some  choice,  more  ex- 
periments are  included  in  the  book  of  laboratory  exercises  than 
can  be  performed  by  all  the  class  within  the  time  ordinarily 
allotted  to  laboratory  work.  Some  of  the  exercises  are  marked 
"Optional,"  while  parts  of  others  are  printed  in  fine  type,  indi- 
cating that  they  may  also  be  omitted  if  time  is  not  available. 


[v] 


FOREWORD  TO  THE  STUDENT 

The  laboratory  is  a  place  for  seeing,  but  it  is  equally  a  place 
for  thinking.  The  laboratory  course  will  include  a  certain 
amount  of  work  that  is  largely  mechanical,  such  as  bending  and 
sealing  glass  tubes,  but  in  the  main  it  is  headwork  rather  than 
handwork  that  counts.  Each  experiment  is  designed  to  bring 
out  an  important  fact  or  to  illustrate  some  principle.  Be  sure 
you  grasp  the  meaning  of  the  experiment  before  you  start  upon 
it.  Study  the  directions  with  the  same  care  that  you  would 
give  to  an  assignment  in  the  textbook  for  recitation. 

Keep  your  desk  neat  and  clean.  A  glance  at  a  student's 
work  desk  will  usually  tell  an  experienced  teacher  exactly  what 
kind  of  work  is  being  done  at  it.  A  dirty  desk,  littered  with 
unnecessary  articles  and  with  apparatus  awkwardly  put  to- 
gether, almost  surely  indicates  a  lack  of  understanding  of  the 
work  in  hand  and  a  willingness  to  accept  without  question  any 
thing  the  book  says.  A  clean  desk  and  well-connected  appara- 
tus indicate  clear  thinking  and  real  joy  in  the  work.  Be 
determined  to  prove  the  truth  of  things  for  yourself. 

Get  ready  for  an  experiment  before  you  start  it.  In  addition 
to  your  burner,  have  at  hand  all  the  materials  called  for  at  the 
beginning  of  the  directions  for  each  experiment.  There  will  be 
four  or  five  bottles  of  the  most-used  reagents  on  your  desk,  and 
the  others  will  be  on  the  shelf  for  general  use. 

Be  cautious  about  reagents.  Some  of  the  substances  you  will 
handle  are  more  or  less  poisonous.  Some  are  inflammable, 
while  others  are  explosive  under  certain  conditions.  Many  at- 
tack the  skin  and  occasion  painful  wounds.  There  are  many 

[  vii  ] 


reagents  in  the  laboratory  that  cannot  be  mixed  with  safety. 
Take  time  to  read  labels  correctly  ;  you  cannot  afford  to  make 
mistakes.  There  is  absolutely  no  danger  if  you  follow  direc- 
tions, but  you  must  exercise  the  same  care  as  does  a  mechanic 
in  a  machine  shop  or  a  driver  of  an  automobile. 

Be  economical.  Reagents  are  expensive,  and  someone  must 
pay  the  bills.  Experiments  nearly  always  go  better  with  small 
quantities  than  with  large  ones,  and  are  more  quickly  carried 
out.  Every  good  chemist  works  with  the  smallest  quantities 
possible  — partly  to  save  money,  but  more  largely  to  save  time 
and  learn  more.  Stick  to  the  quantities  specified  in  the  notes. 
Actually  weigh  or  measure  these  quantities  until  you  can  guess 
very  closely  how  much  a  gram  or  a  cubic  centimeter  really  is. 

Make  your  notes  correspond  to  the  questions.  The  ques- 
tions to  which  written  answers  are  to  be  given  are  indicated  by 
letters  of  the  alphabet  in  parentheses ;  thus,  (a).  Always  place 
the  same  letter  before  your  reply  to  connect  it  with  the  proper 
question.  Note  the  abbreviations  frequently  used  in  the  direc- 
tions: grams  (g.),  centimeters  (cm.),  cubic  centimeters  (cc.), 
liters  (I.).-  The  abbreviation  R.  S.  means  that  a  given  reagent 
is  to  be  sought  on  the  general  reagent  shelf  and  not  upon  your 
own  desk.  Some  exercises  are  marked  "Optional,"  while  por- 
tions of  others  are  printed  in  fine  type  to  indicate  that  these 
portions  are  also  optional.  The  best  students  of  the  class  will 
find  time  to  perform  at  least  some  of  these  optional  experi- 
ments. The  list  of  apparatus  and  materials  required  for  op- 
tional experiments  is  included  in  the  list  at  the  beginning  of 
the  exercise,  but  is  inclosed  in  parentheses. 


[  viii  ] 


CONTENTS 

EXERCISE  PAGE 

1.  The  Metric  System  (see  Appendix).    Methods  of  Measuring 

the  Length,  Volume,  and  Weight  of  Various  Objects     .     .  3 

2.  The   Bunsen  Burner 6 

3.  The  Manipulation  of  Glass  Tubing 8 

4.  Some  Further  Preliminary  Manipulations n 

5.  Chemical  Compounds  ;  Elements  ;   Chemical  Action  (Chemical 

Change);    Chemical   Affinity 14 

6.  The  Collection  of  Gases ;   the  Preparation  of  Oxygen   (Pre- 

liminary)     16 

7.  The  Laboratory  Preparation  of  Oxygen;  Properties  of  Oxygen  17 

8.  A  Study  of  Some  of  the  Changes  taking  Place  when  a  Sub- 

stance Burns 20 

9.  The  Preparation  and  Properties  of  Hydrogen 22 

10.  The  Combustion  of  Hydrogen 26 

11.  Oxidation  and  Reduction 27 

12.  The  Measurement  of  Gas  Volumes .      .  29 

13.  Boiling  Points  and  Freezing  Points ;  Amorphous  and  Crystalline 

Substances 30 

14.  The  Process  of  Distillation;  the  Composition  of  Water  .     .  31 

15.  The  Properties  of  Hydrogen  Peroxide 33 

16.  The  Preparation  and  Properties  of  Nitrogen 34 

17.  The  Molecular  Weight  of  Oxygen;  the  Percentage  of  Oxygen 

in  Potassium  Chlorate  (Optional) 36 

18.  The  Determination  of  the  Formula  of  Copper  Oxide  (Optional)  39 

19.  Some  Further  Manipulations 41 

20.  The  Formation  of  Charcoal  and  Coke 44 

21.  A  Further  Study  of  Carbon 45 

22.  The  Formation  of  Carbon  Dioxide  from  Carbon  and  Oxygen  46 

23.  A  Study  of  Carbon  Dioxide 47 

24.  The  Preparation  and  Properties  of  Carbon  Monoxide  (Optional)  49 

25.  The  Determination  of  the  Relative  Volumes  of  Oxygen  and 

Nitrogen  in  the  Air  (Optional) 51 

[ix] 


.<• 

EXERCISE  PAGE 

26.  A  Study  of  Solutions 53 

27.  The  Determination  of  the  Solubility  of  Common  Salt  ...  55 

28.  The  Preparation  and  Properties  of  Chlorine 56 

29.  The  Preparation  and  Properties  of  Hydrogen  Chloride  and  of 

Hydrochloric  Acid        58 

30.  Sodium ;  Sodium  Hydroxide 59 

31.  The  Properties  of  Acids,  Bases,  and  Salts 60 

32.  The  Ratio  of  Acid  to  Base  in  Neutralization  .....  61 

33.  Carbonic  Acid  and  its  Salts  (Carbonates) 63 

34.  A  Method  for  determining  whether  a  Given  Liquid  is  a  Con- 

ductor of  Electricity  (Optional) 64 

35.  The  Displacement  of  Metals  from  their  Compounds  (the  Dis- 

placement Series) 66 

36.  The  Preparation  and  Properties  of  Ammonia 67 

37.  The  Preparation  and  Properties  of  Nitric  Acid     ....  69 

38.  The  Properties  of  the  Salts  of  Nitric  Acid  (Nitrates)  ...  71 

39.  The  Preparation  and  Properties  of  Some  of  the  Oxides  of 

Nitrogen 72 

40.  The  Properties  and  Chemical  Conduct  of  Sulfur    ....  73 

41.  The  Preparation  and  Properties  of  Hydrogen  Sulfide  ...  74 

42.  The  Preparation  and  Properties  of  the  Salts  of  Hydrosulfuric 

Acid  (Sulfides) 76 

43.  Sulfur  Dioxide  and  Sulfurous  Acid 77 

44.  A  Study  of  Sulfuric  Acid 79 

45.  Salts  of  Sulfuric  Acid  (Sulfates) 80 

46.  The  Preparation  and  Properties  of  Hydrogen  Fluoride    .  81 

47.  The  Preparation  and  Properties  of  Bromine 82 

48.  The  Preparation  and  Properties  of  Iodine 83 

49.  The  Compounds  of  the  Halogens   (Chlorine,  Bromine,  and 

Iodine)  with  Hydrogen 85 

50.  The  Salts  of  the  Binary  Acids  of  the  Halogens:  Chlorides, 

Bromides,  and  Iodides 86 

51.  The  Properties  of  Gasoline  and  Kerosene;  Acetylene  ...  88 

52.  Some  Derivatives  of  Methane  :  Chloroform,  Carbon  Tetra- 

chloride,  lodoform 89 

53.  A  Study  of  Flames 90 

54.  Some  Coal-Tar  Compounds  (Optional) 92 

55.  The  Sugars 94 

56.  The  Composition  of   Milk 95 

57.  The  Determination  of  the  Fat  Present  in  Milk  (Optional)    .  97 

[x] 


58.  A  Study  of   Starch 99 

59.  Textile  Fibers;  Paper 100 

60.  The  Preparation  and  Properties  of  Common  Alcohol  .      .      .  102 

61.  The  Composition  of  Flour 103 

62.  The  Action  of  Preservatives   (Optional) 104 

63.  Acetic  Acid:  a  Study  of  Vinegar ...  106 

64.  Esters  :  Fats  and  Oils 108 

65.  Methods  for  distinguishing  between  Butter  and  Oleomargarine 

(Optional) 109 

66.  Proteins no 

67.  Phosphorus  and  its  Compounds in 

68.  Some  Compounds  of  Arsenic 112 

69.  A  Study  of  Antimony 113 

70.  A  Study  of  Bismuth 114 

71.  Compounds  of  Silicon 115 

72.  Compounds   of   Boron - 115 

73.  Colloids;    Emulsions 117 

74.  General  Methods  for  the  Preparation  of  the  Compounds  of 

the  Metals 119 

75.  The  Compounds  of  Sodium 120 

76.  The  Determination  of  the  Weight  of  Common  Salt  obtained  by 

adding  Hydrochloric  Acid  to  a  Definite  Weight  of  Sodium 

Bicarbonate  (Optional) 121 

77.  Some  Compounds  of  Potassium 123 

78.  The  Properties  of  Ammonium  Compounds 124 

79.  Detection  of  Compounds  of  the  Alkali  Metals  (Optional)    .  125 

80.  The  Preparation  and  Properties  of  Soap     .      .      .      .      .      .  125 

81.  A  Study  of  Some  of  the  Compounds  of  Calcium  .      .      .      .  126 

82.  The  Properties  of  Bleaching  Powder 130 

83.  Hard  Waters  and  Methods  for  Softening  them      .      .      .      .  131 

84.  Testing  the  Acidity  of  Soils 132 

85.  Action  of  Hard  Waters  on  Soap  (Optional) 133 

86.  Magnesium  and  its  Compounds 134 

87.  Zinc  and  its  Compounds 135 

88.  Rubber 136 

89.  Aluminium  and  its  Compounds 137 

90.  A  Study  of  the  Use  of  Aluminium  Sulfate  in  the  Purification 

of  Water 138 

91.  Reactions   of   Baking-Powders 139 

92.  Analysis  of  Baking-Powders   (Optional) 140 

[xi] 


93.  The  Use  of  Mordants  in  Dyeing  (Optional) 141 

94.  A  Study  of  Lakes ;  also  the  Effect  of  using  Different  Mor- 

dants with  the  Same  Dye   (Optional) 143 

95.  The  Detection  of  Dyes  in  Foods 144 

96.  Clay;   Portland  Cement;   Mortar 145 

97.  A  Study  of  Iron  and  its  Compounds 146 

98.  The  Removal  of  Stains 147 

99.  A  Study  of  Copper  and  its  Compounds 149 

100.  A  Study  of  Mercury  and  its  Compounds 150 

101.  A  Study  of  Silver  and  its  Compounds 151 

102.  The  Chemistry  of  Photography 152 

103.  Some  Properties  of  Tin 154 

104.  A  Study  of  Lead  and  Some  of  its  Compounds     ....  154 

105.  Simple  Cells  for  producing  Electric  Currents 155 

106.  Paints 157 

107.  A  Study  of  Some  of  the  Compounds  of  Manganese  .      .      .  158 

108.  A  Study  of  Some  of  the  Compounds  of  Chromium    .      .     .  158 

109.  The  Detection  of  Silver,  Lead,  and  Mercury  when  present  in 

the   Same   Solution    (Optional) .  159 

no.  Borax-Bead  Tests  (Optional) 160 

APPENDIX 

The  Metric  System 161 

Table  of  Solubilities  of  Some  of  the  Compounds  of  the 

Metals 163 

Treatment  in  Case  of  Accident 164 

Information  regarding  Apparatus   and  Chemicals     .      .      .  165 

Apparatus  required  for  Each  Student 166 

Apparatus  to  be  left  on  Each  Desk 168 

Reagents  on  Each  Desk 1 68 

General  Apparatus  for  Ten  Students 168 

Chemicals  on  Reagent  Shelf 169 

Chemicals  required  for  a  Class  of  Ten 171 

Table  of  Constants *75 


[xii] 


LABORATORY  PRACTICE  IN  CHEMISTRY 


EXERCISE  1 


THE    METRIC    SYSTEM    (SEE    APPENDIX).     METHODS    OF 

MEASURING    THE    LENGTH,    VOLUME,    AND    WEIGHT    OF 

VARIOUS  OBJECTS 

Apparatus.  Graduated  test  tube  ;  2  balances,  one  for  weighing  from 
0.5  g.  to  1000  g.  and  the  other  from  o.oi  g.  to  50  g.  or  75  g. ;  watch 
glass ;  loo-cc.  beaker. 

Materials.    5  or  6  g.  common  salt. 

NOTE.  Exercise  i  is  inserted  for  the  benefit  of  any  students  who  may 
not  have  studied  physics  or  who  for  any  reason  may  not  be  familiar 
with  the  metric  system  or  with  the  process  of  weighing  small  objects. 


TUT 


TTTT 


TTTTTnT 


10 


FIG.  i.  Ten-centimeter  scale 

1.  Length.    By  means  of  the  lo-centimeter  scale   (Fig.  i) 
measure    and   (a)  record    the    length    of    various    pieces    of 
apparatus  included  in  your  outfit,  as  a  test  tube, 

a  file,  and  a  blowpipe,  (b)  What  is  the  diameter 
of  your  filter  paper?  Estimate  the  lengths  of 
various  objects,  as  a  pencil,  a  test  tube;  then 
measure.  Continue  until  you  can  approximate 
the  lengths  of  such  objects. 

2.  Volume.    By  means  of  a  graduated  test  tube 
or  cylinder  (Fig.  2)  measure  in  cubic  centimeters 
and  (c)  record  the  volumes  of  various  test  tubes, 
beakers,  and  flasks  included  in  your  outfit.    In 
reading  off  the  amount  of  the  liquid  in  a  grad- 
uated tube,  always  read  from  the  lower  part  of 

the  meniscus ;  that  is,  the  curved  surface  of  the  liquid  (Fig.  3). 

3.  Weight.    The  platform  balance  (Fig.  4)   is  adapted   for 
weighing  objects  that  weigh  from  0.5  g.  to  1000  g.    It  is  not  so 

[3] 


FIG.  2.  Gradu- 
ated test  tube 


\ 


FIG.  3.    Taking  the  reading  of  the 

volume  of  a  liquid  in  a  graduated 

test  tube 


sensitive  as  the  chemical  balance  (Fig.  3)  or  the  hornpan  bal- 
ance (Fig.  6),  either  of  which  is  sensitive  to  o.oi  g.  These 
smaller  balances,  however,  while  very  sensitive,  can  only  be  used 
in  weighing  objects  up  to  50  g. 
or  75  g.  at  the  most.  The  arms 
of  the  chemical  balance  (Fig.  5) 
can  be  raised  or  lowered  by 
turning  the  screw  D.  When  the 
balance  is  not  in  use  the  arms 
should  always  be  lowered  so 
that  the  pans  rest  on  the  top  of 
the  base  of  the  balance. 

Before  using  a  balance  the 
two  arms  must  be  in  exact  equilibrium.  In  the  case  of  the 
platform  balance,  if  the  arms  are  in  equilibrium,  then  the 
pointer  A  (Fig.  4)  will  point  directly  to  the  zero  mark  on 
the  scale  B  when  the  pans  are  at  rest ;  if  it  does  not,  bring  it  to 
this  position  by  adjusting  the  screw  C.  In  the  more  sensitive 
balance  (Fig.  5)  the 
pointer  A  must 
swing  equal  dis- 
tances on  either  side 
of  the  zero  point 
marked  on  the  scale 
B  when  the  pans 
move  up  and  down; 
if  it  does  not,  make 
it  do  so  by  turning 
the  screw  C. 
When  the  balance 


J 

FIG.  4. 


Platform  balance  for  weighing  objects  up 
to  1000  g. 


is  adjusted,  the  object  to  be  weighed  is  placed  on  one  pan 
(the  left  is  more  convenient)  and  weights  (Figs.  7  and  8)  are 
added  to  the  other  until  the  arms  are  in  equilibrium.  In 
the  case  of  the  platform  balance  the  equivalent  of  10  g.  or 
any  fraction  of  this  weight  is  obtained  by  simply  moving  the 

[4] 


,c 


FIG.  5.    Chemical  balance  for  accurately 
weighing  small  objects  up  to  50  g.  or  75  g. 


slide  D  (Fig.  4)  along  the  scale.  When  the  arms  are  in  equilib- 
rium the  weights  are  carefully  noted  and  recorded.  After 
a  little  practice  one  can  ap- 
proximate closely  the  weight 
of  the  object  and  thus  know 
about  what  weights  must  be 
added.  Care  must  be  taken 
to  read  and  record  the 
weights  accurately.  It  is  a 
good  plan  to  check  the  read- 
ings of  the  weights  by  first 
noting  the  vacant  places  in 
the  box  in  which  the  weights 
are  kept. 

It  is  not  advisable  here  to 
give  detailed  directions  in 
regard  to  the  use  of  the  balance;  such  information  will  be 
furnished  by  the  instructor.  It  must  be  remembered,  however, 
that  a  balance  is  a  delicate  instrument 
and. must  be  given  the  same  care  as  a 
watch.  The  objects  weighed  must  be 
clean  and  dry,  and  the  weights  must 
be  handled  only  with  small  forceps 
(Fig.  9).  Under  no  condition  must 
chemicals  be  placed  directly  on  the 
scalepan.  Such  substances  as  sugar 
and  salt  may  be  weighed  out  on  paper ; 
others,  which  act  upon  paper,  must  be 
weighed  in  a  small  beaker  or  on  a 
watch  glass. 

After  studying  the  construction  of 
the  balance  and  the  method  of  using 
it,  weigh  and  (d)  record  the  weights  of 
various  small  objects,  as  a  porcelain  crucible,  a  watch  glass. 

(e)  What  is  the  approximate  weight  of  a  nickel  five-cent  piece  ? 

[5] 


FIG.  6.  Hornpan  balance. 
This  can  be  used  in  place 
of  the  chemical  balance,  but 
it  is  not  so  convenient, 
although  less  expensive 


Accurately  balance  a  watch  glass  on  the  scalepan  (Fig.  5), 
using  either  the  weights  or,  better,  a  small  pill  box  and  fine 
sand,  and  weigh 
out  on  this  ex- 
actly 5.2  g.  of 
common  salt. 

Clean  and  dry 
a  small  beaker 
and  balance  it  on 


FIG.  7.  Weights  used  with 
the  chemical  balance 


FIG.  8.  Weights  used  with 
the  platform  balance 


the  scalepan ;  then  remove  it  from  the  pan  and  pour  into  it,  as 
nearly  as  possible,   locc.  of  distilled 
water  ( measured  in  the  graduated  tube 
included  in  your  laboratory  outfit) .  Re- 
weigh,  and  (/)  note  the  weight  of  the 
water,  (g)  How  do  you  account  for  the    FlG  Q  The  proper  way  of 
fact  that  it  does  not  weigh  exactly  10  g.  ?  handling  weights 


EXERCISE  2 
THE  BUNSEN  BURNER 

Apparatus.   A  Bunsen  burner  and  rubber  tube ;  wing-top  burner. 
Materials.   Wooden  splints  (i  cm.  x  12  cm.  is  a  convenient  size); 
white  cardboard  (locm.  square). 

1.  Construction.  The  Bunsen  burner  is 
a  form  of  apparatus  used  for  burning  gas 
and  is  the  common  burner  employed  in  the 
laboratory.  It  consists  of  the  tube  A  (Fig. 
10),  screwed  into  the  base  C.  The  tube  has 
two  small  openings  near  its  lower  part.  A 
small  band  B,  provided  with  similar  open- 
ings, fits  around  the  lower  part  of  the  tube 
in  such  a  way  that  the  openings  of  the  tube 
may  be  closed  or  kept  open  by  turning  the 
band.  Gas  is  admitted  through  D  by  means  of  rubber  tubing 
connected  with  the  gas  pipe. 

[6] 


FIG.  10.    A   Bunsen 
burner 


Unscrew  the  tube  and  examine  the  different  parts  of  the 
burner  (Fig.  n) ;  then  put  them  together  again.  Turn  on  the 
gas  and  at  once  ignite  it  by  holding  a  lighted 
match  4  or  5  cm.  above  the  top  of  the  tube. 
The  supply  of  gas  should  be  adjusted  so  as 
to  give  a  flame  about  10  cm.  high.  The  gas 
flowing  through  the  tube  mixes  with  the  air 
drawn  in  through  the  openings  in  the  lower 
part  of  the  tube  and  burns  with  an  almost 
nonluminous  flame.  If  the  band  is  adjusted 
so  as  to  close  the  openings,  the  flame  becomes 
luminous.  Always  use  the  nonluminous 
flame  unless  otherwise  directed. 

Turn  off  the  gas,  slip  over  the  end  of  the 
tube  the  so-called  " wing-top"  burner   (A, 
Fig.  12),  and  again  ignite  the  gas.    Notice 
that  the  flame  is  spread  out.   This  form  of  the  flame  is  al- 
ways used  in  bending  glass  tubes,  as  is  described  in  Exercise  3. 


FIG.  ii.    The  parts 
of  a  Bunsen  burner 


FIG.  12.  Wing-top  burner 
(A),  and  also  a  Bunsen 
burner  fitted  with  a  wing- 
top  burner 


FIG.  13.  Testing  the  relative  tem- 
peratures of  the  different  parts  of 
a  flame  by  holding  a  cardboard 
in  the  flame 


2.  Determining  the  relative  temperatures  of  the  flame.    Hold 
a  small  wooden  splint  horizontally  in  the  base  of  the  Bunsen 

[7] 


flame  for  two  or  three  seconds  and  (a)  note  the  results.  In  the 
same  way  determine  the  relative  temperatures  of  various  parts 
of  the  flame.  Turn  the  gas  down  until  the  flame  is  7  or  8  cm. 
in  height ;  then  quickly  thrust  a  piece  of  white  cardboard  verti- 
cally through  the  center  of  the  flame,  the  lower  end  of  the  card- 
board resting  against  the  top  of  the  burner  (Fig.  13).  Remove 
the  cardboard  before  it  is  ignited,  and  from  the  scorched 
portions  note  the  relative  temperatures  of  different  parts  of  the 
flame,  (b)  Draw  a  diagram  to  illustrate  your  results. 


B 


f 


EXERCISE  3 

THE  MANIPULATION  OF  GLASS  TUBING 

Apparatus.  Burner  and  wing-top  attachment;  hard-glass  test  tube 
with  cork  to  fit ;  cork-borers ;  triangular  file ;  round  file. 

Materials.  Glass  tubing  (soft),  external  diameter  6  mm.;  about 
30  cm.  of  glass  rod. 

1.  To  fit  a  tube  with  stopper  and  glass  tube  as  shown  in 
Fig.  14.  In  all  operations  requiring  the  application  of  a  strong 
heat  to  glass,  the  heat  must  be  applied  gently 
at  first ;  otherwise  a  strain  is  produced  which  is 
apt  to  break  the  glass.  For  the  same  reason 
highly  heated  glass  must  be  cooled  slowly. 

From  one  of  the  lengths  of  soft-glass  tubing 
cut  a  piece  about  1 5  cm.  in  length.  To  do  this, 
place  the  tubing  on  the  desk  and,  holding  it 
firmly  with  your  hand,  draw  the  edge  of  a  tri- 
angular file  across  the  point  at  which  you  wish 
to  cut  the  glass.  After  the  glass  has  been 
scratched,  take  the  tube  in  the  hands  with  the 
thumbs  placed  near  together  just  back  of  the 
scratch  (Fig.  15)  and  gently  pull  the  glass  apart, 
at  the  same  time  exerting  a  slight  pressure  with 
the  thumbs.  If  the  tube  does  not  yield  readily  to  a  gentle 
pressure,  a  deeper  scratch  must  be  made.  In  the  case  of  large 


FIG.  14.   A  test 

tube  fitted  with 

cork  and  glass 

tube 


FIG.  15.  After  the  glass  has  been 

scratched  with  a  file,  it  is  pulled 

apart  as  shown 


tubing  it  may  be  found  necessary  to  file  a  groove  around  the 
glass.  The  edges  of  the  cut  tube  will  be  sharp  and  should  be 
rounded  by  being  rotated  in  the  tip  of  the  Bunsen  flame 
(Fig.  16). 

To  bend  the  glass  tubing,  first 
heat  it,  at  the  point  where  you 
wish  to  bend  it,  in  the  luminous 
Bunsen  flame  spread  out  by  means 
of  the  wing-top  burner  (Fig.  17). 
Hold  the  tube  lengthwise  in  the 
flame,  gently  rotating  it  so  that  all 
sides  may  be  equally  heated.  Con- 
tinue the  heating  until  the  glass  bends  easily,  then  remove  it 
from  the  flame  and  quickly  bend  it  to  a  right  angle  B  (Fig.  14). 
Great  care  must  be  taken  to  heat  the  tube  uniformly;  other- 
wise the  bore  of  the  tube  will  be  contracted  (A,  B,  Fig.  18), 
forming  a  bend  which  is  not  only 
unsightly  but  is  easily  broken. 

Next  select  a  good  cork  of  such 
a  size  that  the  smaller  end  will  just 
enter  the  hard-glass  test  tube.  Soften 
the  cork  by  rolling  it  between  the 
desk  and  a  block  of  wood.  The  next 
step  is  to  insert  the  glass  tube,  pre- 
pared as  directed  above,  into  and 
through  the  cork.  To  do  this,  select 
a  borer  (A,  Fig.  19)  slightly  smaller 
than  the  tube.  Place  the  cork  on  a 
piece  of  cardboard  on  the  desk  and, 
holding  the  rod  B  firmly  with  your 
hand,  cut  through  the  cork  with  the 
borer,  not  by  punching  but  by  rotating  the  borer  under  gentle 
pressure  (Fig.  19).  Care  must  be  taken  to  keep  the  borer  at  a 
right  angle  to  the  top  and  base  of  the  cork.  The  hole  should  be 
straight  and  smooth. 

[9] 


FIG.  1 6.   Rounding  the  sharp 
ends  of  a  glass  tube  or  rod 


The  end  of  the  glass  tube  is  now  moistened  with  water  or 
vaseline  and  gently  pushed  through  the  cork  by  a  screwlike 
motion.  If  the  hole  is  too 
small  to  admit  the  tube 
when  a  gentle  pressure  is 
applied,  it  must  be  slightly 
enlarged  with  a  round  file. 
It  must  be  remembered 
that  the  glass  is  fragile  and 
will  break  if  more  than  a 
gentle  pressure  is  applied;  and  if  broken  while  grasped  in  the 
hand,  a  serious  cut  may  result.  Now  put  the  cork  in  the 
test  tube  and  set  the  apparatus  aside  for  use  in  preparing 
oxygen  (Exercise  7). 

B 


FIG.  17.    Heating  tube  before  bending  it 


B 

FIG.  18.     Poorly    bent 
tubes 


FIG.  19.  Boring  a  hole  through 
a  cork  stopper 


2.  To  make  glass  stirring-rods.  These  should  be  about 
15  cm.  in  length.  Cut  the  glass  rod  into  pieces  of  this  length^ 
using  the  method  employed  in  cutting  glass  tubing.  Round  the 
ends  of  the  rods  by  heating  in  the  Bunsen  flame  (Fig.  16). 
Place  the  finished  rods  in  the  desk  for  future  use. 


[10] 


EXERCISE  4 

SOME  FURTHER  PRELIMINARY  MANIPULATIONS 

Apparatus.  2  beakers  ;  graduated  tube ;  glass  stirring-rod ;  funnel ; 
evaporating-dish ;  mortar  and  pestle. 

Materials.   Filter  paper ;  about  i  g.  each  of  sand  and  salt. 

NOTE.  Since  a  Bunsen  burner  and  a  ring  stand  are  on  each  desk, 
they  will  not  be  included  in  the  list  of  apparatus  in  subsequent  exercises. 

1.  Pouring  liquids  from  one  vessel  to  another.  In  doing  this 
care  must  be  taken  to  prevent  the  liquid  from  running  down 
the  side  of  the  vessel  from  which  it  is  poured.  A  glass  rod 


FIG.  20.  Pouring  liquid  from  a  bottle 
into  a  tube 


FIG.  21.    Removing  the 
stopper  from  a  bottle 


(prepared  in  Exercise  3)  should  be  held  lightly  against  the 
rim  of  the  vessel,  as  shown  in  Fig.  23.  The  liquid  will  flow 
down  the  rod. 

Fill  a  beaker  with  water  and  transfer  it  slowly  to  another 
vessel  without  using  the  glass  rod ;  repeat,  using  the  glass  rod. 

In  pouring  liquids  from  bottles  a  glass  rod  may  be  used, 
or  the  neck  of  the  bottle  may  be  placed  lightly  against  tKe 
rim  of  the  vessel  into  which  the  liquid  is  being  poured  (Fig.  20). 
This  will  prevent  the  liquid  from  running  down  the  side  of  the 

[111 


FIG.  22.  The  proper  way  of  folding  a  filter 
paper  and  fitting  it  in  a  funnel 


bottle.  The  stopper  must  never  be  laid  on  the  desk.  Catch  it 
between  the  fingers,  as  shown  in  Fig.  21,  leaving  the  hand  free 
to  grasp  the  bottle,  as 
shown  in  Fig.  20. 

2.  Decantation;  fil- 
tration. It  is  often  nec- 
essary to  separate  a 
liquid  from  a  finely 
divided  solid  which  is 
suspended  in  it.  This 
may  be  done  by  one 
of  the  two  following 
methods : 

When  the  solid  is 
heavy  and  readily  settles  to  the  bottom  of  the  vessel,  the 
liquid  may  be  carefully  poured  off,  or  decanted.  Sand  sus- 
pended in  water  may  be  separated  from  the  liquid  in  this  way. 

Usually,  however,  the  solid  will 
not  readily  settle  or  will  do  so  only 
after  long  standing.  In  such  cases 
the  mixture  is  filtered,  by  pouring  it 
on  a  filter  paper,  which  allows  the 
liquid  to  run  through  but  retains 
the  solid.  The  method  of  folding  the 
filter  paper  and  fitting  it  into  a  fun- 
nel is  shown  in  Fig.  22.  The  paper 
(A,  Fig.  22)  is  first  folded  along  a 
diameter  into  halves  (B) ;  then,  at 
right  angles  to  the  first  fold,  into 
quarters  (C).  The  folded  filter  is 
then  opened  so  as  to  form  a  cone 
(D),  half  of  which  is  composed  of 
three  thicknesses  of  paper,  and  the  remainder  of  one  thickness. 
Now  fit  the  cone  into  a  funnel  of  such  a  size  that  the  paper  does 
not  quite  reach  the  top  (E).  The  paper  must  fit  the  funnel 

[12] 


FIG.  23.  Separating  a  solid  from 
a  liquid  by  filtration 


accurately ;  if  it  does  not,  make  it  do  so  by  varying  the  fold. 
Place  the  paper  in  the  funnel  and  thoroughly  wet  it  with  water. 
After  the  water  has  run  through,  press  the  paper  firmly  against 
the  sides  of  the  funnel  with  the  finger  so  as  to  remove  any  air 
bubbles  between  the  paper  and  the  glass.  The  filter  is  now 
ready  for  use  (Fig.  23).  The  process  of  filtration  not  only 
enables  us  to  separate  liquids  from 
solids  but  also  certain  solids  from 
each  other,  as  is  illustrated  in  the 
following  experiment: 

3.  Separating  two  solids  from  each 
other.  Weigh  out  about  i  g.  each  of 
sand  and  common  salt  and  mix  them 
thoroughly  by  grinding  them  to- 
gether in  a  mortar.  Transfer  the  mix- 
ture to  a  small  beaker,  add  10  cc.  of 
water,  and  stir  the  mixture  for  two 
or  three  minutes  with  a  glass  rod. 
Filter  off  the  undissolved  solid,  col- 
lecting the  filtrate  (that  is,  the  liquid 


FIG.  24.  Evaporating  a  liquid 


that  runs  through  the  paper)  in  a  beaker  (Fig.  23).  Pour  the 
filtrate  into  an  evaporating-dish.  Support  the  dish  on  a  ring 
stand  (Fig.  24)  and  evaporate  the  liquid  by  a  gentle  heat.  With- 
draw the  heat  as  soon  as  the  water  is  evaporated,  (a)  What  is 
the  solid  left  on  the  filter  paper?  (b)  What  is  the  solid  left  in 
the  evaporating-dish  (taste)  ?  (c)  Can  all  mixtures  of  solids  be 
separated  by  the  above  method? 


[13 


EXERCISE  5 


CHEMICAL  COMPOUNDS ;  ELEMENTS ;  CHEMICAL  ACTION 
(CHEMICAL  CHANGE) ;  CHEMICAL  AFFINITY 

Apparatus.    Test  tube  (ordinary);  hard-glass  test  tube. 
Materials.    Iron  wire  10  cm.  in  length ;  wooden  splints ;   2  or  3  g. 
sugar  ;  i  g.  mercuric  oxide. 

Caution.  In  heating  any  substance  in  a  tube  care  must  al- 
ways be  taken  never  to  point  the  tube  toward  yourself  or  any- 
one else;  for  in  some  cases  the  hot  contents  may  be  ejected 
from  the  tube  and,  if  they  come  in  contact  with  the  body}  may 
cause  more  or  less  serious  injury.  The 
correct  way  of  heating  a  tube  is  shown 
in  Fig.  25. 

1.  Hold  a  piece  of  iron  wire  in  the 
Bunsen  flame  for  a  few  seconds,   (a)  Is 
the  iron  changed?    Examine  the  wire 
after  it  has  cooled,   (b)  Have  the  prop- 
erties of  the  iron  been  changed  by  the 
heating?    (c)  Has  the  iron  undergone 
a  chemical  change? 

Repeat  the  experiment,  using  a 
wooden  splint,  (d)  Does  the  wood 
undergo  a  chemical  change?  (e)  Con- 
trast the  action  of  heat  on  the  iron  and 
the  wood. 

2.  Introduce  2  or  3  g.  of  sugar  into  a  clean,  dry  test  tube. 
This  is  done  by  placing  the  sugar  near  the  end  of  a  narrow 
strip  of  folded  paper  and  introducing  it  into  the  tube  held  in  a 
horizontal  position,  as  shown  in  Fig.  26.    If  the  paper  is  now 
turned  over,  the  sugar  is  deposited  in  the  bottom  of  the  tube ; 
the  paper  is  then  carefully  withdrawn,  leaving  the  sides  of  the 
tube  perfectly  clean. 

[14] 


FIG.  25.  The  correct  way  of 
heating  a  liquid  in  a  tube 


Fold  a  piece  of  paper  12  or  15  cm.  in  length  into  the  form  of 
a  band  from  i  to  2  cm.  in  width,  place  the  band  about  the  upper 
end  of  the  tube,  and  grasp  the  ends  of  the  band  between  your 
thumb  and  forefinger,  as  shown  in  Fig.  25.  The  tube  may  now 
be  heated  without  burning  your  fingers.  Heat  the  sugar  gently 
in  the  tip  of  the  flame,  as  shown  in  Fig.  25,  moving  the  tube 
about  in  the  flame  so  as  to  heat  the  sugar  uniformly.  (/)  Note 
carefully  the  changes  that  take  place. 
Continue  the  heating  as  long  as  any 
changes  occur  and  until  a  black  solid 
is  left.  This  solid  is  the  element 

FIG.  26.    Introducing   a   small 
carbon,    (g)  Contrast  the  properties      amount  of  solid  into  a  tube 

(taste,  color,  solubility  in  water)  of 

the  sugar  and  the  carbon,  (h)  What  is  the  source  of  the 
carbon?  (i)  What  grounds  have  you  for  assuming  that  a 
chemical  change  has  taken  place? 

3.  Introduce  (Fig.  26)  about  i  g.  of  mercuric  oxide  into  a 
clean,  dry,  hard-glass  test  tube.    Heat  the  oxide  strongly,  in- 
serting a  glowing  splint  from  time  to  time  down  into  the  tube. 
Note  that  the  splint  bursts  into  a  flame.    This  is  caused  by  the 
invisible  gas  oxygen,  which  is  evolved  when  the  mercuric  oxide 
is  heated.    Note  the  residue  clinging  in  droplets  on  the  sides 
of  the  tube  near  the  bottom.   This  is  the  element  mercury 
(quicksilver). 

4.  (;)  Define  the  following  terms,  illustrating  each  by  ex- 
amples taken  from  the  above  experiments:  element,  chemical 
compound,  chemical  change,  and  chemical  affinity. 


[15] 


EXERCISE  6 


THE   COLLECTION  OF  GASES;   THE  PREPARATION   OF 
OXYGEN   (PRELIMINARY) 

Apparatus.  Three  250-00.  wide-mouthed  bottles  ;  3  pieces  of  window 
glass  10  cm.  square ;  pneumatic  trough ;  rubber  tubing ;  2  test  tubes. 

Materials.  4  g.  potassium  chlorate  ;  i  g.  powdered  manganese  diox- 
ide ;  wooden  splints. 

1.  Collection  of  gases.  Fill  a  250-00.  wide-mouthed  bottle 
with  water.  Cover  its  mouth  with  a  glass  plate,  being  careful 
to  exclude  all  air  bubbles.  Hold  the  plate  firmly  in  place, 
invert  the  bottle, 
and  dip  its  mouth 
into  the  water  in  a 
pneumatic  trough. 
Remove  the  glass 
plate,  (a)  Why 
does  the  water  re- 
main in  the  bot- 
tle? Now,  without 
changing  the  posi- 
tion of  the  bottle, 
fill  it  with  exhaled 
air  by  placing  one 
end  of  a  piece  of 
glass  or  rubber 
tubing  under  the  FlG  27  Filling  a  bottle  with  exhaled  air 

mouth  of  the  bot- 
tle and  blowing  gently  through  the  other  end  (Fig.  27). 

Before  the  bottle,  so  filled,  is  removed  from  the  trough, 
cover  its  mouth  tightly  with  a  glass  plate.  The  bottle  so 
covered  may  then  be  placed  on  the  desk  either  right  side  up  or 
in  an  inverted  position,  according  to  the  weight  of  the  gas. 

[161 


Fill  a  bottle  with  exhaled  air  and  then  transfer  the  air  to 
another  bottle,  (b)  Draw  a  diagram  to  show  a  good  method  of 
doing  this. 

2.  Preparation  of  oxygen  from  mercuric  oxide.    Recall  3, 
Exercise  5. 

3.  Preparation  of  oxygen  from  potassium  chlorate  (prelimi- 
nary experiment).    Select  two  test  tubes  of  the  same  size  and 
clean  and  dry  them  thoroughly.    Into  the  one  introduce  2  g.  of 
potassium  chlorate ;  into  the  other  introduce  2  g.  of  potassium 
chlorate  mixed  intimately  with  i  g.  of  manganese  dioxide. 

Now  heat  the  contents  of  the  two  tubes  gently,  applying  the 
flame  so  that  both  tubes  are  equally  heated.  Repeatedly  thrust 
a  glowing  splint  into  each  tube  in  order  to  detect  any  oxygen 
that  may  be  evolved,  (c)  Note  the  results,  (d)  What  effect 
has  the  manganese  dioxide  ?  (e)  From  which  tube  is  the  greater 
amount  of  oxygen  evolved  (consult  text)  ? 

EXERCISE  7 

THE  LABORATORY  PREPARATION  OF  OXYGEN ; 
PROPERTIES  OF  OXYGEN 

NOTE.  In  some  exercises  certain  experiments  are  optional.  The 
paragraphs  describing  the  optional  experiments  are  in  fine  print.  The 
apparatus  and  materials  required  for  the  optional  experiments  are  given 
in  the  list,  but  the  names  are  inclosed  in  parentheses  ( — ). 

Apparatus.  Hard-glass  test  tube  fitted  with  cork  and  delivery  tube 
as  shown  in  Fig.  28  (use  apparatus  constructed  in  Exercise  3,  Fig.  14); 
four  250-cc.  wide-mouthed  bottles;  4  pieces  of  glass  10  cm.  square; 
pneumatic  trough;  deflagrating-spoon ;  (evaporating-dish ;  funnel). 

Materials.  6  g.  potassium  chlorate  ;  3  g.  manganese  dioxide  ;  wooden 
splints ;  i  g.  of  sulfur ;  picture-frame  wire  20  cm.  long ;  bit  of  cotton ; 
(filter  paper). 

1.  Usual  laboratory  method  for  preparing  oxygen.  Mix  in- 
timately on  paper  6  g.  of  potassium  chlorate  and  3  g.  of  man- 
ganese dioxide.  The  presence  of  impurities  in  the  materials 

[17] 


may  lead  to  a  serious  explosion  when  heat  is  applied;  hence 
test  a  small  portion  of  the  mixture,  say  0.5  g.,  by  heating  it 
in  a  test  tube.  In  the  absence  of  impurities  the  oxygen  is 
evolved  quietly,  unaccompanied  by  any  very  marked  sparking 
in  the  materials. 

If  pure,  transfer  the  remainder  of  the  mixture  to  a  hard- 
glass  tube,  insert  the  cork,  and  arrange  the  apparatus  as  shown 


FIG.  28.   Apparatus  required  for  preparing  oxygen  by  heating  a  mixture  of 
potassium  chlorate  and  manganese  dioxide 

in  Fig.  28.  The  bottles  must  be  completely  filled  with  water 
and  inverted  in  the  water  in  the  pneumatic  trough.  Now  hold- 
ing the  burner  in  the  hand,  heat  the  mixture  gently  with  a  small 
flame,  applying  the  heat  at  first  to  the  upper  part  of  the  mix- 
ture. The  flame  must  not  strike  the  upper  part  of  the  test 
tube,  as  the  cork  may  be  burned.  At  first  the  heat  expands  the 
air,  and  a  few  bubbles  of  air  escape ;  then  the  oxygen  is  evolved. 
Regulate  the  heat  so  as  to  secure  a  uniform  and  not  too  rapid 
evolution  of  the  gas.  Collect  three  or  four  2  50-0:.  wide-mouthed 
bottles  of  the  gas.  Before  the  heat  is  withdrawn,  remove  the 
cork  from  the  tube,  (a)  What  is  the  reason  for  this  precau- 

[18] 


rfi 


tion  ?  Slip  a  piece  of  glass  over  the  mouth  of  each  bottle ;  then 
remove  the  bottle  from  the  water  and  set  it  in  an  upright  posi- 
tion on  your  desk,  keeping  the  mouth  of  the  bottle  covered  by 
the  glass. 

(b)  What  is  the  source  of  the  oxygen?  (c)  What  is  the  func- 
tion of  the  manganese  dioxide?  Place  the  tube  and  contents 
aside  for  use  in  3,  below. 

2.  Properties  and  chemical  conduct  of  oxygen. 
(d)  Note  the  properties  (color,  odor,  taste,  solu- 
bility in  water)  of  the  gas.    (The  slight  cloud 
that  is  often  present  when  oxygen  is  prepared 
from  potassium  chlorate  is  due  to  an  impurity 
and  will  disappear  if  the  gas  is  allowed  to  stand 
over  water.) 

(e)  Repeatedly  thrust  a  glowing  splint  into 
a  bottle  of  the  gas  and  note  the  results. 

Heat  some  sulfur  in  a  deflagrating-spoon  un- 
til it  begins  to  burn.    (/)  Note  the  color  and 
size  of  the  sulfur  flame,    (g)  Now  lower  the  burning  sulfur 
into  a  bottle  of  oxygen  (Fig.  29)  and  note  the  change. 

Tip  the  piece  of  picture-frame  wire  with  sulfur  by  wrapping 
a  small  bit  of  cotton  about  the  end  of  the  wire  and  dipping 
this  into  melted  sulfur  (for  this  purpose  melt  a  little  sulfur  in 
a  deflagrating-spoon).  Ignite  the  sulfur  by  holding  it  in  a 
Bunsen  flame  for  an  instant,  and  then  thrust  the  end  of  the 
wire  into  a  bottle  of  oxygen,  (h)  Record  the  results. 

(/)  Summarize  the  properties  and  chemical  conduct  of  oxygen. 

3.  Separation  of  the  compounds  present  in  the  residue  left  in  the 
preparation  of  oxygen.     Heat  the  tube  containing  the  residue  left  in 
the  preparation  of  oxygen  until  no  more  oxygen  is  evolved.    (;')  What 
two  compounds  are  present  in  the  residue  (consult  text)  ?    One  of  these 
is  soluble  in  water  and  the  other  insoluble.    Devise  a  method  for  the 
separation  of  the  two  compounds  and  carry  out  the  process,    (k)  De- 
scribe the  method  and  record  your  results. 


FIG.  29.  Burning 
sulfur  in  oxygen 


[19 


EXERCISE  8 


A  STUDY  OF  SOME  OF  THE  CHANGES  TAKING  PLACE 
WHEN  A  SUBSTANCE  BURNS 

Apparatus.  Porcelain  crucible ;  pipestem  triangle  to  support  the  cru- 
cible while  being  heated ;  small  beaker ;  glass  tube  12  or  15  cm.  in  length ; 
25o-cc.  wide-mouthed  bottle;  balance;  (pneumatic  trough;  large  wide- 
mouthed  bottle  or  large  beaker). 

Materials.  2  or  3  g.  finely  powdered  iron  (use  iron  reduced  by  alco- 
hol); candle;  10  cc.  limewater  (R.S.)  (the  abbreviation  R.  S.  sig- 
nifies that  the  reagent  is  found  on  the  reagent  shelf) ;  (pellet  of 
phosphorus  size  of  a  pea,  to  be  obtained  from  instructor  when  needed ; 
iron  wire  about  15  cm.  long). 

1.  The  burning  of  metals.  Iron  may  be  taken  as  a  typical 
metal,  since  the  metals  in  general  burn  in  the  same  way  as  iron. 
Place  2  or  3  g.  of  powdered 
iron  in  a  porcelain  crucible 
and  weigh  to  o.oi  g.,  (a)  re- 
cording the  weight.  Now  heat 
the  crucible  (Fig.  30)  until  the 
iron  begins  to  glow  (burn) ; 
then  withdraw  the  heat,  (b) 
Does  the  iron  continue  to 
burn?  When  the  crucible  is 
cool,  reweigh.  (c)  Compare 
the  weight  of  the  product  with 
that  of  the  unburned  iron  and 


account    for    any    change    in 
weight. 

2.  The  burning  of  a  candle. 
Pour  5  cc.  of  clear  limewater 

FIG.  30.  Burning  iron  powder 

into  a  small  beaker,  and  by 

means  of  a  glass  tube  gently  blow  exhaled  air  through  the 
liquid  (Fig.  31).    Note  that  the  liquid  soon  becomes  cloudy, 

[20] 


or  "milky,"  in  appearance.   This  change  is  due  to  the  action 
of  the  colorless  and  odorless  gas  known  as  carbon  dioxide  (which 
is  present  in  the  exhaled  air)  and  con- 
stitutes a  good  test  for  this  gas. 

Hold  a  cold,  dry,  wide-mouthed 
bottle  over  a  candle  flame  as  shown  in 
Fig.  32.  Note  the  film  of  moisture 
(water)  collecting  on  the  bottle,  (d) 
What  is  the  source  of  this  water  ?  After 
one  or  two  minutes  remove  the  bottle 
quickly  and  pour  into  it  5  cc.  of  clear 
limewater.  Place  the  palm  of  the  hand 
tightly  over  the  mouth  of  the  bottle 
and  shake  the  contents,  (e)  Note  any 
change  in  appearance  of  the  limewater. 
(/)  What  conclusions  do  you  draw  from 
the  experiment?  (g)  How  do  you  ac- 
count for  the  fact  that  a  burning  candle 
gradually  disappears,  while  iron,  on 
burning,  increases  in  weight  ? 


Limewater 


FIG.  31.   Blowing  exhaled 
air  through  limewater 


3.  The  burning  of  phosphorus.  (Precaution-)  Phosphorus  must  be 
kept  and  handled  only  under  water  ;  otherwise  it  may 
ignite  and  serious  results  follow.  Cover  the  bottom 
of  a  pneumatic  trough  with  water  to  a  depth  of  2 
or  3  cm.  On  the  water  float  a  porcelain  crucible. 
Ask  the  instructor  to  place  in  this  crucible  a  small 
piece  of  phosphorus  the  size  of  a  pea. 

Ignite  the  phosphorus  by  touching  it  with  the  end 
of  a  hot  wire  or  file  (Fig.  33),  and  quickly  invert 
over  the  crucible  a  wide-mouthed  bottle  or  beaker, 
being  careful  to  keep  the  mouth  of  the  bottle  below 
the  surface  of  the  water.  The  white  powder  formed 
by  the  burning  phosphorus  floats  in  the  air  in  the 
bottle  but  is  gradually  dissolved  in  the  water. 

Leave  the  bottle  in  position  until  the  powder  has 
entirely  disappeared.      Note  that  the  water  has  risen  in  the  bottle. 
(h)  How  do  you  account  for  this  fact?    (i)  Is  your  conclusion  in 

[21] 


FIG.  32.  Collect- 
ing the  hot  gases 
evolved  by  a 
burning  candle 


accord  with  the  results  obtained  by  burning  iron  in  air?  (;')  Suppose 
it  were  possible  for  you  to  collect  and  weigh  the  white  powder  formed 
by  the  burning  phosphorus,  how  should  you  expect  its  weight  to  com- 
pare with  that  of  the  original  phosphorus? 


FIG.  33.  Burning  phosphorus  in  air 

Remove  the  crucible  from  the  water  and  place  it  on  a  ring  stand  in 
the  hood  (Fig.  30).  Now  heat  the  crucible  very  gently  (care)  so  as 
to  ignite  any  unburned  phosphorus.  When  cool,  the  crucible  may  then 
be  cleaned  and  put  away. 


EXERCISE  9 

THE  PREPARATION  AND  PROPERTIES  OF  HYDROGEN 

Apparatus.  Test  tube ;  apparatus,  bottles,  and  trough  as  shown  in 
Fig.  35  (the  bottles  are  25o-cc.  wide-mouthed);  beaker;  stirring-rod; 
6o-cc.  bottle;  forceps;  4  glass  plates;  (evaporating-dish ;  funnel). 

Materials.  Bit  of  sodium  (size  of  a  small  pea) ;  wooden  splints  ;  10  g. 
mossy  zinc  ;  i  cc.  copper  sulfate  solution  (R.  S.)  ;  i  cc.  kerosene  (R.S.)  ; 
sulf uric  acid ;  (filter  paper). 

1.  Preparation  from  water.  Fill  a  test  tube  with  water  and 
invert  it  in  a  beaker  of  water.  Wrap  a  piece  of  sodium  in  a  bit 
of  filter  paper  previously  moistened  with  kerosene.  The  sodium, 
when  wrapped  in  the  paper,  must  be  small  enough  to  enter  the 
tube  easily.  Raise  the  inverted  test  tube  until  its  mouth  dips 
just  below  the  surface  of  the  water  in  the  beaker  and,  grasping 

[22] 


the  sodium  and  paper  with  the  forceps,  quickly  insert  the  so- 
dium into  the  test  tube  (Fig.  34).  Stand  at  arm's  length,  as  a 
slight  explosion  sometimes  occurs.  Notice  that  the  sodium  de- 
composes the  water,  liberating  a  gas  which  is  caught  in  the  tube. 
When  the  action  has  ceased,  place  your  thumb  tightly  over  the 
mouth  of  the  tube  to  prevent  the  gas  from  escaping,  and  bring 
the  tube  to  an 
upright  position. 
After  lighting  a 
splint,  remove  the 
thumb  from  the 
tube  and  quickly 
bring  the  flame  to 
the  mouth  of  the 
tube,  (a)  Does  the 
gas  act  like  oxy- 
gen? (b)  What  is 
the  source  of  the 
gas?  (c)  Record 
any  other  methods 
that  may  be  em- 
ployed for  obtaining  it  from  the  same  source  (consult  text). 
2.  Preparation  from  acids  (usual  laboratory  method).  Pre- 
pare a  hydrogen  generator  according  to  Fig.  35.  D  represents 
a  wide-mouthed  bottle  of  about  250-0:.  capacity.  The  gas- 
delivery  tube  B,  C  is  the  same  as  that  used  in  the  preparation  of 
oxygen  (Fig.  28).  A  rubber  stopper  for  the  bottle  D  is  pref- 
erable, although  a  good  cork  stopper  will  do  provided  it  can 
be  made  air-tight.  The  funnel  tube  A  must  extend  nearly  to  the 
bottom  of  the  bottle.  Put  10  g.  of  the  mossy  zinc  into  the  bottle 
D.  Pure  sulfuric  acid  will  not  react  readily  with  pure  zinc; 
hence  it  is  advisable  to  add  to  the  zinc  from  8  to  10  drops  of  a 
solution  of  copper  sulfate,  which  will  start  the  reaction.  Now 
pour  through  the  funnel  tube  into  the  bottle  just  enough  water 
to  cover  the  zinc. 

[23] 


FIG.  34.   Preparing  hydrogen  by  the  action  of  sodium 
on  water 


Prove  that  the  apparatus  is  air-tight  by  blowing  into  the  de- 
livery tube  C  until  the  water  is  forced  nearly  to  the  top  of  the 
funnel  tube ;  then  quickly  close  the  rubber  tube  either  by  tightly 
pinching  it  or  by  placing  the  tongue  firmly  against  its  end.  If  the 
apparatus  is  air-tight,  the  water  in  the  funnel  tube  will  not  fall. 

Now  prepare  some  dilute  sulfuric  acid  by  pouring  slowly,  a 
few  drops  at  a  time,  15  cc.  of  concentrated  acid  into  a  beaker 
containing  50  cc.  of  water.  Stir  the  water  with  a  glass  rod  while 


FIG.  35.  Apparatus  used  for  preparing  hydrogen  by  the  action  of  sulfuric 

acid  on  zinc 

the  acid  is  being  added.  Notice  that  the  acid  is  poured  into 
the  water  —  never  the  reverse.  (Sulfuric  acid  mixes  with  water, 
evolving  a  great  deal  of  heat.  Unless  the  above  precautions  are 
taken,  the  heat  may  be  so  intense  as  to  cause  the  liquid  to  spat- 
ter, and  the  hot  acid  in  contact  with  the  skin  causes  a  serious 
burn.) 

Cool  the  dilute  acid  and  pour  a  few  drops  of  it  through  the 
funnel  tube  (A).  Hydrogen  is  evolved  as  soon  as  the  acid  comes 
in  contact  with  the  zinc.  Enough  of  the  acid  must  be  added 
from  time  to  time  to  cause  a  gentle  and  continuous  evolution 
of  the  gas.  An  excess  of  the  acid  should  be  avoided,  however, 
or  the  action  will  become  too  violent,  and  a  large  quantity  of 
zinc  will  have  to  be  added  at  the  close  of  the  exercise. 

[24] 


It  is  evident  that  the  gas  which  passes  over  first,  is  a  mixture 
of  hydrogen  and  air.  The  student  must  remember  that  such  a 
confined  mixture  of  hydrogen  and  air  explodes  with  great  vio- 
lence if  ignited ;  hence  see  that  the  end  of  the  delivery  tube  is 
not  brought  near  any  flame.  Determine  when  the  hydrogen  is 
free  from  air  by  repeatedly  collecting  a  test  tube  full  of  gas 
and  igniting  it,  holding  the  tube  mouth  downward,  (d)  Why 
downward?  If  pure,  the  gas  burns 
quietly;  otherwise  there  is  a  slight 
explosion.  After  all  the  air  has  been 
expelled  from  the  generator,  collect 
four  bottles  (250-0:.  wide-mouthed) 
of  the  gas ;  then  cover  the  mouth  of 
each  bottle  with  a  glass  plate  and 
place  the  bottles  mouth  downward 
on  your  desk. 

(e)  What  is  the  source  of  the  hydro- 
gen ?  (/)  What  is  the  use  of  the  zinc  ? 
(g)  Why  is  mossy  zinc  used? 

Remove  the  cork  from  the  gener- 
ator, add  a  few  more  pieces  of  zinc, 
and  set  aside  for  use  in  4,  below. 

3.  Properties  and  chemical  conduct  of  hydrogen.  Thrust  a 
lighted  splint  into  a  bottle  of  the  gas  held  mouth  downward 
(Fig.  36).  Slowly  withdraw  the  splint  and  again  thrust  it  into 
the  gas.  (h)  Describe  the  results,  (i)  What  do  they  prove? 

Fill  a  small  (6o-cc.)  wide-mouthed  bottle  or  test  tube  one 
third  full  of  water  and  invert  it  in  a  pneumatic  trough.  Displace 
the  remaining  water  with  hydrogen  from  one  of  the  bottles. 
(;)  What  does  the  bottle  now  contain?  Withdraw  it  from  the 
water  and,  holding  it  at  arm's  length,  quickly  bring  it,  mouth 
downward,  over  a  flame,  (k)  What  do  the  results  prove? 

Uncover  a  bottle  (mouth  upward)  of  the  gas.  After  a  minute, 
test  for  the  presence  of  hydrogen  with  a  lighted  splint.  Repeat, 
holding  the  bottle  mouth  downward.  (I)  What  do  the  results 

[25] 


FIG.  36.    Thrusting  a  lighted 
splint  into  hydrogen 


show?    (m)   Summarize  briefly  the  properties  and  chemical 
conduct  of  hydrogen. 

4.  Without  removing  the  fragments  of  any  undissolved  zinc,  pour  the 
liquid  left  in  the  generator  D  in  experiment  2  into  an  evaporating-dish 
and  boil  gently  on  a  ring-stand  support.  As  soon  as  white  crusts  begin 
to  form  on  the  side  of  the  dish,  just  above  the  liquid,  filter  the  hot  liquid 
into  a  beaker  and  set  the  filtrate  aside  to  cool,  (n)  How  does  the  prod- 
uct which  separates  from  the  filtrate  compare  in  properties  with  the 
original  zinc?  (0)  What  is  the  product  (consult  text)?  (p}  How  does 
it  differ  in  composition  from  sulfuric  acid  ? 

EXERCISE  10 
THE  COMBUSTION  OF  HYDROGEN 

Apparatus.  Hydrogen  generator  A  (Fig.  37)  attached  by  rubber 
tubing  E  to  a  drying-tube  B.  This  tube  is  filled  with  granulated  cal- 
cium chloride,  held  in  place  by  loose  plugs  of  cotton  placed  at  each  end 
of  the  tube.  C  is  a  glass  tube  drawn  to  a  jet ;  rubber  tube  C  (Fig.  35). 

Materials.  Granulated  calcium  chloride  sufficient  to  fill  tube  B 
(Fig.  37);  cotton  sufficient  to  use  in  tube  B ;  8  g.  zinc;  i  cc.  copper 
sulfate  solution  (R.S.);  dilute  sulfuric  acid  for  preparing  hydrogen  (see 
Exercise  9) ;  picture-frame  wire  10  cm.  long ;  bit  of  charcoal. 

1.  Charge  the  generator  A  (Fig.  37)  with  6  or  8  g.  of  zinc, 
add  the  solution  of  copper  sulfate,  cover  with  water,  and  add 
dilute  sulfuric  acid  as  in  2,  Exercise  9,  taking  all  the  precautions 
noted  in  the  exercise.  The  hydrogen  passes  from  the  bottle 
through  the  tube  E,  is  freed  from  all  moisture  in  passing  over 
the  calcium  chloride  in  the  tube  B,  and  escapes  through  the 
tube  C.  Slip  a  piece  of  rubber  tubing  over  the  tube  C,  collect 
samples  of  the  gas  in  test  tubes  over  water,  and  test  with  a  flame 
to  see  whether  the  hydrogen  is  free  from  air.  After  all  the  air 
has  been  expelled,  wrap  a  towel  carefully  about  the  generator, 
remove  the  rubber  tubing,  and  cautiously  ignite  the  hydrogen 
escaping  at  the  end  of  tube  C.  The  flame  is  nearly  invisible  and 
is  very  hot.  Test  the  heat  of  the  flame  by  holding  in  it  different 
objects,  such  as  a  splint,  a  piece  of  picture-frame  wire,  a  bit 
of  charcoal. 

[261 


Finally,  hold  over  the  flame  a  cold,  dry  beaker  or  bottle  and 
note  the  liquid  (water)  deposited  on  the  sides  of  the  vessel. 

(a)  Summarize  briefly  all  the  facts  brought  out  in  the  above 
experiment. 


FIG.  37.    Apparatus  for  preparing  and  burning  hydrogen 


EXERCISE  11 

OXIDATION  AND  REDUCTION 

Apparatus.  Hydrogen  generator  and  tubes  as  shown  in  Fig.  38  (^4  is 
the  hydrogen  generator,  B  is  a  drying-tube  filled  with  calcium  chloride, 
C  is  a  straight  glass  tube,  and  D  is  a  hard-glass  test  tube) ;  apparatus 
used  in  preparing  oxygen  (Fig.  28). 

Materials.  2  g.  copper  oxide ;  calcium  chloride  sufficient  to  fill  the 
drying-tube  B ;  8  g.  zinc ;  dilute  sulfuric  acid  for  preparing  hydrogen 
(see  Exercise  9) ;  copper  sulfate  solution  (R.  S.) ;  4  g.  potassium  chlo- 
rate ;  2  g.  manganese  dioxide. 

1.  Remove  the  test  tube  D  (Fig.  38),  introduce  into  it  i  or 
2  g.  of  copper  oxide,  and  set  it  aside. 

[271 


Now  generate  hydrogen  in  A  as  in  Exercise  10.  After  all  the 
air  has  been  expelled  from  the  apparatus  (test  samples  to  be 
certain),  wrap  the  generator  in  a  towel  as  described  in  Exer- 
cise 10.  Then  slip  the  test  tube  D  over  the  exit  tube  C,  as  shown 
in  Fig.  38,  and  cautiously  heat  the  copper  oxide  in  D  to  redness, 
being  careful  to  keep  the  flame  away  from  the  mouth  of  the 
tube.  Note  the  condensation  of  moisture  in  the  cold  portions  of 


FIG.  38.  Reducing  copper  oxide  by  hydrogen 

the  tube,  (a)  Is  there  any  visible  evidence  of  change  in  the 
copper  oxide?  (b)  Explain,  (c)  What  is  the  object  of  the 
calcium  chloride  in  tube  B  ? 

2.  Disconnect,  the  bottle  A  at  ,E  from  the  rest  of  the  ap- 
paratus and  blow  air  through  the  tube  B  to  remove  the  hydro- 
gen present.   Now  connect  the  apparatus  used  in  preparation 
of  oxygen  (Fig.  28)  at  £  to  the  tubes  B,  C,  and  D.    Generate 
oxygen  as  in  Exercise  7  and  conduct  a  slow  current  of  the  gas 
through  B,  C,  and  D,  and  at  the  same  time  gently  heat  the 
residue  in  D.    (d)  Record  your  results. 

3.  (e)  Define  the  terms  reduction  and  oxidation  and  give  an 
example  of  each  from  the  above  experiment. 

[28] 


EXERCISE  12 


THE  MEASUREMENT  OF  GAS  VOLUMES 

Apparatus.    Graduated  tube   and  cylinder  as   shown  in  Fig.   39 ; 
thermometer. 

1.  (It  is  suggested  that  the  instructor  arrange  one  or  more 
pieces  of  apparatus  as  shown  in  Fig.  39.  The  students  will 
then  take  the  readings  and  solve  the 
problems.) 

Partially  fill  a  graduated  tube  with 
water  and  invert  it  in  a  cylinder  (or 
other  vessel)  of  water  as  shown  in  Fig.  39. 
Let  the  water  stand  until  its  temperature 
is  the  same  as  that  of  the  air.  Adjust 
the  tube  until  the  level  of  the  liquid  in- 
side and  outside  of  the  tube  is  the  same ; 
then  take  the  reading  of  the  volume  of 
the  air  in  the  tube.  Note  the  temperature 
of  the  air  (place  a  thermometer  by  the 
side  of  the  tube)  and  likewise  the  pres- 
sure of  the  atmosphere  as  indicated  by  . 

r  ,       FIG.  39.  Measuring  a  vol- 

the  barometer.     Insert  these  values  in      ume  Of  gas  over  water 

their  appropriate  places  in  the  following 

table;  then  calculate  the  volume  which  the  air  would  occupy 

under  standard  conditions. 

Volume  of  air  in  tube cc. 

Temperature  of  air 

Barometric  pressure mm. 

Vapor  pressure  of  water  at  temperature  of  air   (see 

Appendix)       mm. 

Effective  pressure  on  air  in  tube  (barometric  pressure 

less  vapor  pressure) mm. 

Volume  which  the  air  would  occupy  under  standard 

conditions   (calculate) cc. 

[291 


EXERCISE  13 


BOILING  POINTS  AND  FREEZING  POINTS ;  AMORPHOUS 
AND  CRYSTALLINE  SUBSTANCES 

Apparatus.  Flask  (250-00.)  fitted  with  a  two-hole  cork,  through 
which  pass  the  thermometer  B  and  the  glass  tube  C,  as  shown  in 
Fig.  40 ;  test  tube  ;  large  beaker  ;  small  beaker ;  stirring-rod ;  watch  glass. 

Materials.  300  g.  ice ;  100  g.  common  salt ;  3  g.  powdered  alum ;  5  g. 
powdered  sulfur ;  i  cc.  chloroform. 

1.  Boiling  points.    Fill  the  flask  A  (Fig.  40)  about  one  third 
full  of  water.   Lower  the  thermometer  until  the  bulb  is  im- 
mersed in  the  liquid.   Heat  the  water  to  boiling  and  note  the 
temperature  after  it  has  become  constant. 

Apply  a  greater  heat  and  note  the  effect 
upon  the  temperature. 

Raise  the  thermometer  until  the  bulb 
no  longer  touches  the  boiling  water  but  is 
simply  exposed  to  the  steam.  Again  note 
the  temperature,  (a)  Record  your  general 
conclusions. 

2.  Freezing  points.    Place  some  pieces 
of  ice  in  a  small  beaker.   Introduce  15  or 
20  cc.  of  water  and  stir  the  mixture  gently 
with  the  thermometer,   (b)  Continue  until 
the  temperature  is  constant,  and  record  the 
temperature. 

Half  fill  a  test  tube  with  pure  water  and 
dip  the  tube  into  a  freezing-mixture  made 
by  mixing  3  parts  of  powdered  ice  with 
i  part  of  common  salt.  Stir  the  water  in  the  test  tube  from 
time  to  time  with  a  thermometer  and  note  the  temperature  at 
which  it  begins  to  freeze,  (c)  Compare  the  freezing  point  of 
water  with  the  melting  point  of  ice. 

[30] 


FIG.  40.  Apparatus  for 

determining  the  boiling 

point  of  a  liquid 


3.  Add  powdered  sulfur  to  a  test  tube  until  the  tube  is  about 
one  third  filled;  then  heat  the  tube  gently  until  the  sulfur  is 
melted.    Set  the  tube  aside  until  the  liquid  has  again  solidified. 
(d)  Is  the  solid  in  a  crystalline  form? 

4.  Place  about  3  g.  of  powdered  alum  in  a  small  beaker  or 
evaporating-dish  and  dissolve  the  solid  in  as  little  water  as  pos- 
sible, at  room  temperature.    Hasten  the  solution  by  stirring  the 
mixture  with  a  glass  rod.    In  the  same  way  make  a  solution  of 
common  salt.    Set  the  solutions  aside  (do  not  cover  the  beaker 
or  dish)  until  the  next  laboratory  period,    (e)  Note  the  form 
in  which  each  solid  is  deposited,  illustrating,  if  possible,  by  a 
drawing.  (/)  Mention  two  common  methods  used  for  obtaining 
substances  in  crystalline  state. 

5.  Pour  about  i  cc.  of  chloroform  on  a  watch  glass  and  blow 
across  the  surface  of  the  liquid,     (g)  Account  for  its  rapid 
evaporation,    (h)  Note  and  account  for  any  change  in  the  tem- 
perature of  the  glass. 

EXERCISE  14 

THE  PROCESS  OF  DISTILLATION ;  THE  COMPOSITION 
OF  WATER 

Apparatus.  Flask  (250-00.);  Liebig  condenser  (B)  and  connections 
as  shown  in  Fig.  41,  or  apparatus  shown  in  Fig.  42 ;  wire  gauze  12  cm. 
square  ;  watch  glass  ;  glass  rod. 

Materials.  10  cc.  alcohol. 

1.  Distillation  of  water.  (Two  or  more  students  may  work 
together  if  the  apparatus  is  not  available  for  each.)  Connect 
a  Liebig  condenser  B  with  a  250-0:.  flask  A,  as  represented  in 
Fig.  41.  The  flask  is  set  on  a  piece  of  wire  gauze  supported  by 
the  iron  ring  attached  to  the  ring  stand.  The  tube  C  is  con- 
nected with  the  water  pipe  by  means  of  rubber  tubing,  and  a 
current  of  cold  water  is  conducted  through  the  outer  tube  of 
the  condenser,  escaping  through  the  tube  D.  (a)  Why  is  cold 
water  admitted  at  C  rather  than  at  £)? 

[31] 


FIG.  4L   Distilling  water  in  the  laboratory 


Fill  the  flask  one  fourth  'full  of  hydrant  or  well  water  and 
boil  the  water  until  25  cc.  or  more  of  liquid  has  collected  in  the 
receiver  E.  If  con- 
densers are  not  avail- 
able, the  apparatus 
shown  in  Fig.  42  may 
be  used.  The  steam 
from  the  boiling 
water  in  flask  A  is 
condensed  by  con- 
ducting it  through  B 
into  the  test  tube  C, 
which  is  kept  cold  by 
ice  water  in  a  beaker 
or  bottle  D. 

(b)  Compare  the  distillate  (distilled  water)  with  the  hydrant 
water  in  appearance  and  in 
taste. 

Place  4  or  5  drops  of  the 
distilled  water  on  a  watch 
glass  and  evaporate,  hold- 
ing the  watch  glass  10  or 
1 5  cm.  above  the  tip  of  the 
flame,  (c)  Is  there  any 
residue?  (d)  Repeat,  using 
hydrant  water,  (e)  Why  is 
distilled  water  used  in  the 
laboratory  ? 

Repeat  the  distillation, 
using  a  sample  of  muddy 
water  in  A,  and  (/)  record 
the  results. 

2.  Fractional  distillation.  Dip  the  end  of  a  glass  rod  into 
alcohol  and  at  once  bring  it  into  contact  with  a  flame,  (g)  Is  the 
alcohol  inflammable?  Now  distill  (Fig.  41)  a  mixture  of  10  cc. 

[32] 


B 


D 


FIG.  42.  A  simple  apparatus  for  distilling 
small  volumes  of  water 


of  alcohol  (boiling  point  78.3°)  and  30  cc.  of  water.  Collect 
the  first  i  or  2  cc.  of  the  distillate  in  an  evaporating-dish  and 
test  with  a  flame.  In  the  same  way  test  successive  portions  of 
the  distillate,  (h)  Does  there  seem  to  be  a  partial  separation 
of  the  two  liquids  ?  In  this  way  a  mixture  of  liquids  boiling  at 
different  temperatures  may  generally  be  separated  more  or  less 
perfectly.  The  process  is  termed  fractional  distillation. 

3.  The  composition  of  water.  Recall  the  experiments  in 
Exercise  10,  also  in  Exercise  n.  (i)  What  do  these  experiments 
show  in  regard  to  the  elements  present  in  water?  (;)  Can  you 
suggest  any  modification  of  Exercise  n  that  would  enable  you 
to  determine  the  relative  weights  of  the  elements  present  in 
water  ? 

EXERCISE  15 

THE  PROPERTIES  OF  HYDROGEN  PEROXIDE 

Apparatus.    3  test  tubes ;  funnel ;  small  beaker. 

Materials.  10  cc.  hydrogen  peroxide  solution;  i  g.  manganese  diox- 
ide ;  wooden  splint ;  filter  paper ;  i  cc.  of  starch  solution  (R.  S.) ;  small 
crystal  of  potassium  iodide. 

1.  (a)  What  is  the  strength  of  the  hydrogen  peroxide  solution 
sold  by  druggists  (consult  text)  ?    Pour  3  cc.  of  the  solution 
into  a  large  test  tube  and  add  i  or  2  g.  of  finely  powered  man- 
ganese dioxide,    (b)  Test  the  gas  evolved  with  a  glowing  splint. 

Filter  the  mixture  remaining  in  the  tube.  The  solid  is  the 
unchanged  manganese  dioxide,  (c)  In  what  other  experiment 
has  manganese  dioxide  brought  about  a  change  without  appar- 
ently undergoing  any  change  itself? 

2.  To  i  cc.  of  starch  solution  add  4  cc.  of  water.    Dissolve 
in  this  a  small  crystal  of  potassium  iodide  and  then  add  a  few 
drops  of  a  solution  of  hydrogen  peroxide.    The  peroxide  liber- 
ates the  iodine  from  the  potassium  iodide,  and  the  resulting 
free  iodine  colors  the  starch  paste,    (d)  Note  the  color. 


33] 


EXERCISE  16 

THE  PREPARATION  AND  PROPERTIES  OF  NITROGEN 

Apparatus.  Apparatus  shown  in  Fig.  43  :  A  is  the  hard-glass  tube 
included  in  your  outfit  and  is  supported  on  a  ring  stand.  It  contains  a 
spiral  of  fine  copper  gauze  (from  60  to  70  mesh),  made  by  rolling  up  a 
piece  of  gauze  about  10  x  12  cm.  The  gauze  must  fit  the  tube  loosely 
so  that  it  can  be  pushed  easily  to  the  middle  of  the  tube.  B  is  a  i -liter 
(or  larger)  bottle  (the  ordinary  5-pint  acid  bottle  serves  the  purpose 
well).  The  glass  tube  D  extends  to  the  bottom  of  the  bottle  B  and  is 
connected  to  a  water  tap  by  means  of  a  rubber  tube ;  2  wide-mouthed  bot- 
tles (250-cc.);  pneumatic  trough ;  wing-top  burner  (hydrogen  generator). 

Materials.  Pieces  of  copper  gauze  (from  60 to 70  mesh)  10  x  12  cm.; 
splint;  (zinc  and  dilute  sulfuric  acid  for  generating  hydrogen). 

1.  Preparation  of  nitrogen  from  the  air.  When  air  is  passed 
through  a  tube  containing  hot  copper,  the  oxygen  in  the  air 


FIG.  43.  Preparing  nitrogen  by  passing  air  over  hot  copper 

combines  with  the  copper,  leaving  the  nitrogen.  This  gives  a 
simple  method  for  preparing  nitrogen.  Proceed  as  follows: 
Arrange  the  apparatus  as  shown  in  Fig.  43  and  see  that  the  parts 
are  all  well  connected.  The  bottle  B  is  empty  at  the  beginning, 
while  the  wide-mouthed  bottles  inverted  in  G  are  completely 

[34] 


filled  with  water.  Now  heat  the  copper  spiral  in  the  tube,  using 
a  wing-top  burner,  until  the  gauze  is  dull  red;  this  will  re^ 
quire  about  five  minutes.  Then  continue  the  heating,  and  at 
the  same  time  force  a  slow  current  of  water  into  B  at  a  rate 
not  to  exceed  1 1.  in  twenty  minutes.  The  water  entering  B 
forces  the  air  in  the  bottle  over  the  hot  copper,  which  unites 
with  the  oxygen,  while  the  nitrogen  passes  on  and  is  collected 
in  the  bottles  in  G.  When  two  of  these  bottles  are  filled,  dis- 
connect the  apparatus  and  withdraw  the  heat,  (a)  Why  dis- 
connect the  apparatus  before  withdrawing  the  heat? 

Remove  the  bottles  from  G  as  in  the  preparation  of  oxygen 
and  hydrogen,  keeping  each  covered  with  a  glass  plate.  Since 
the  copper  removes  only  the  oxygen,  the  nitrogen  prepared  by 
this  method  will  contain  small  percentages  of  the  rare  elements 
present  in  the  atmosphere.  Since  they  are  present  in  such  small 
amounts  (about  i  per  cent  of  the  original  air)  and  resemble 
nitrogen  so  closely,  they  do  not  materially  affect  the  properties 
of  the  nitrogen. 

2.  Properties  and  chemical  conduct  of  nitrogen,    (b)  Note 
the  color,  odor,  taste,  and  solubility  in  water  of  nitrogen.  Insert 
a  burning  splint  into  a  bottle  of  the  gas.    (c}  Is  nitrogen  com- 
bustible?   (d)  Is  it  a  supporter  of  combustion?  (e)  Summarize 
briefly  the  properties  of  nitrogen. 

3.  Since  the  copper  spiral  is  partly  oxidized  to  copper  oxide  in  the 
experiment,  and  since  you  will  want  the  copper  spiral  in  a  later  experi- 
ment, it  is  advisable,  if  time  is  available,  to  reduce  the  spiral  back  to 
copper.    To  do  this,  connect  your  hydrogen  generator  to  the  tube  A  con- 
taining the  spiral,  and  after  all  the  air  is  expelled  from  the  apparatus 
(test  as  in  Exercise  9)  gently  heat  the  tube  containing  the  gauze,  while 
passing  a  slow  current  of  hydrogen  through  the  tube.    The  gauze  is  soon 
reduced,  as  shown  by  the  change  in  color.    Leave  the  copper  in  the  tube 
until  cold.    (/)  Why  until  cold  ? 


[35] 


EXERCISE  17 


THE  MOLECULAR  WEIGHT  OF  OXYGEN  ;  THE  PERCENT- 
AGE OF  OXYGEN  IN  POTASSIUM  CHLORATE  (OPTIONAL) 

Apparatus.  Apparatus  shown  in  Fig.  44  (A,  B,  C,  is  the  apparatus 
used  in  the  preparation  of  oxygen,  Fig.  28.  Care  must  be  taken  that 
the  glass  tube  and  cork  fit  the  test  tube  air-tight,  so  that  any  gas  gen- 
erated in  the  tube  can  escape  only  through  the  exit  tube  C.  D  is  an 
ordinary  5-pint  acid  bottle,  which  must  be  completely  filled  with  water 
and  inverted  in  the  water  in  the  trough);  cork  to  fit  the  bottle  D; 
porcelain  crucible;  iron  wire  10  or  12  cm.  long;  graduated  cylinder. 

Materials.  3  g.  powdered  manganese  dioxide ;  7  g.  potassium  chlorate. 

1.  Principle  involved.  The  molecular  weight  of  a  gaseous 
element  or  compound  is  determined  by  obtaining  the  weight  of 
a  definite  volume  of  the  gas.  Since  the  direct  weighing  of 
a  gas  is  not  easily 
carried  out,  it  is 
better  to  adopt 
some  method  which 
avoids  the  direct 
weighing  of  the  gas. 
In  the  case  of  oxy- 
gen this  is  readily 
done  by  selecting  a 
solid  compound  of 
oxygen  (such  as  po- 


FIG.  44.   Determining  the  volume  of  a  definite  weight 
of  oxygen 


tassium  chlorate) 
which  will  evolve 
the  gas  when  heated.  The  difference  in  the  weights  of  the  com- 
pound before  and  after  heating  gives  the  weight  of  the  oxygen 
evolved.  By  collecting  the  oxygen  evolved  it  is  easy  to  deter- 
mine its  volume.  We  have  then  the  volume  occupied  by  a 
definite  weight  of  oxygen,  and  from  these  results  we  can  cal- 
culate the  weight  of  22.4!.  of  the  gas.  Proceed  as  follows: 

[36] 


2.  Preparation  of  the  oxygen.  In  your  porcelain  crucible 
place  about  3  g.  of  finely  powdered  manganese  dioxide  and  heat 
it  strongly  in  the  Bunsen  flame  for  about  five  minutes,  stirring 
it  from  time  to  time  with  a  clean  iron  wire  or  end  of  file ;  then 
withdraw  the  flame  and  allow  the  powder  to  cool. 

Thoroughly  clean  and  dry  your  hard-glass  test  tube  (A,  Fig. 
44)  and  weigh  it  accurately  to  o.oi  g.  Record  the  weight  in  the 
table  below.  Weigh  out  roughly  between  6  and  7  g.  of  potas- 
sium chlorate  and  introduce  this  into  the  tube  (Fig.  26),  taking 
care  that  none  of  the  chlorate  adheres  to  the  sides  of  the  tube. 
Accurately  weigh  (to  o.oi  g.)  the  tube  and  contents  as  before, 
recording  the  weight  in  the  table.  Now  introduce  into  the 
tube  the  manganese  dioxide  contained  in  the  crucible,  and  again 
weigh  to  o.oi  g.  and  record  the  weight  in  the  table.  Mix  the 
two  compounds  in  the  bottom  of  the  tube  by  gently  rotating 
the  tube  and  by  tapping  it  against  your  hand.  Connect  the 
tube  as  shown  in  Fig.  44  and  arrange  the  exit  tube  so  that  all 
gas  evolved  will  be  collected  in  the  bottle  Z),  which  must  be 
completely  filled  with  water.  Now  hold  the  burner  in  your 
hand  and  heat  the  mixture  gently.  Oxygen  is  evolved  and  is 
collected  in  the  bottle.  Regulate  the  flame  so  that  the  bubbles 
of  the  gas  can  easily  be  counted.  Continue  the  heating  until 
the  gas  is  no  longer  evolved.  Remove  the  cork  from  the  test 
tube  and  carefully  set  the  tube  aside  to  cool. 

Now  adjust  the  bottle  containing  the  oxygen  so  that  the 
level  of  the  water  inside  and  outside  the  bottle  is  the  same 
(it  may  be  necessary  to  add  more  water  to  the  trough) ;  then, 
while  holding  it  in  this  position,  push  a  cork  tightly  into  the 
bottle  and  quickly  bring  the  bottle  to  an  upright  position  on 
the  desk.  Next,  take  the  temperature  of  the  water  in  the 
trough  and  the  reading  of  the  barometer  and  record  the  figures 
in  the  table.  If  time  is  not  available  the  bottle  containing  the 
oxygen  (and  some  water),  and  the  tube  A,  in  which  the  oxygen 
was  prepared,  can  be  set  aside  and  the  experiment  concluded 
later. 

[37] 


3.  Determination  of  the  weight  of  22.4  1.  of  oxygen.   Weigh 
to  o.o i  g.  the  test  tube  and  residue  left  after  preparing  the  oxy- 
gen, recording  the  weights  in  the  table.    The  loss  in  the  weight 
of  the  tube  and  contents  on  heating  equals  the  weight  of  the 
oxygen  evolved.   Now  measure  the  volume  of  the  oxygen  in  the 
bottle  by  adding  water  from  the  graduated  cylinder  and  record 
the  volume.   The  volume  occupied  by  the  oxygen  is  under 
laboratory  conditions.    Reduce  this  to  standard  conditions  and 
record  the  -result  in  the  table.    You  now  have  the  volume  under 
standard  conditions  occupied  by  a  known  weight  of  oxygen. 
From  these  results  calculate  the  weight  of  1 1.  of  oxygen  under 
standard  conditions  and  from  this,  the  weight  of  22.4!.  of 
oxygen.    Record  the  weights  in  the  table. 

4.  Calculation  of  percentage  of  oxygen  in  potassium  chlorate. 
All  the  oxygen  evolved  in  the  above  experiment  comes  from 
the  potassium  chlorate ;  moreover,  if  the  experiment  is  carried 
out  as  directed,  all  the  oxygen  in  the  potassium  chlorate  is 
evolved,  since  the  mixture  is  heated  until  the  gas  ceases  to  be 
given  off.    From  the  weight  of  the  potassium  chlorate  taken 
and  the  weight  of  the  oxygen  evolved,  calculate  (see  table)  the 
percentage  of  oxygen  in  the  potassium  chlorate  and  record  your 
results  in  the  table. 

Weight  of  test  tube  A g. 

Weight  of  test  tube  plus  potassium  chlorate    ...  g. 

Weight  of  potassium  chlorate  taken  (equals  the  dif- 
ference -of  the  two  weights  given  above)   ...  g. 

Weight  of  the  test  tube  plus  potassium  chlorate  plus 

manganese   dioxide g. 

Weight  of  the  test  tube  and  residue  left  after  heating  g. 

Weight  of  oxygen  evolved  (equals  loss  in  weight  of 

tube  and  contents  on  heating) g. 

Temperature  of  water  in  pneumatic  trough    .     .     . 

Reading  of  barometer mm. 

Volume  of  oxygen  in  bottle  D cc. 

Volume  of  this  oxygen  under  standard  conditions  (cal- 
culate)   cc. 

[38] 


Weight  of  i  1.  of  oxygen  under  standard  conditions 

(calculate) g. 

Weight  of  22.4  1  of  oxygen  under  standard  conditions 
(calculate).  (The  number  expressing  the  weight 
of  22.4 1.  of  oxygen  is  the  same  as  the  number  ex- 
pressing the  molecular  weight  of  the  gas.)  .  .  g. 

Percentage  of  oxygen  in  potassium  chlorate  (calculate 
from  the  following  formula) : 

Weight  of  oxygen  evolved  X  100       _  J  percentage  of  oxygen  in 
Weight  of  potassium  chlorate  taken  ~~  \     potassium  chlorate      % 

AVERAGE  OF 

RESULTS 

RESULTS  OBTAINED 

THEORETICAL        OBTAINED          BY  CLASS 

Molecular  weight  of  oxygen  .      .       32 
Percentage  of  oxygen  in  potas- 
sium chlorate 39.2 

(a)  What  are  the  sources  of  error  in  the  above  experiment? 

EXERCISE  18 

THE  DETERMINATION  OF  THE  FORMULA  OF  COPPER 
OXIDE  (OPTIONAL) 

Apparatus.  Apparatus  shown  in  Fig.  45  :  A  is  the  hard-glass  tube 
used  in  Exercise  16,  B  is  a  hydrogen  generator,  and  C  is  a  25o-cc.  wide- 
mouthed  bottle  filled  with  water  and  inverted  in  a  pneumatic  trough. 
Porcelain  boat  about  8  cm.  long  and  6  or  7  mm.  wide. 

Materials,  i  or  2  g.  black  finely  powdered  copper  oxide ;  zinc  and 
dilute  sulfuric  acid  for  preparing  hydrogen  as  in  Exercise  10. 

1.  Thoroughly  clean  and  dry  the  porcelain  boat  and  weigh 
it  to  o.o i  g.,  recording  the  weight  in  the  table  below.  Nearly 
fill  the  boat  with  the  copper  oxide  and  again  weigh  to  o.oi  g. 
and  record  the  weight.  Carefully  introduce  the  boat  and 
copper  oxide  into  the  glass  tube  A  (Fig.  45)  and  connect  the  ap- 
paratus (air-tight)  as  shown  in  the  figure.  Now  generate  hydro- 
gen in  bottle  B  as  directed  in  Exercise  10.  When  the  apparatus 
is  filled  with  hydrogen  (note  all  precautions  given  in  Exercise 
10),  carefully  heat  the  porcelain  boat  in  the  tube  with  a  wing- 
top  burner,  maintaining  all  the  time  a  slow  current  of  hydrogen 

[391 


through  the  tube.  Gradually  increase  the  heat  and  continue  the 
operation  for  about  twenty  minutes,  collecting  any  excess 
hydrogen  in  the  bottle  C.  The  hydrogen  reduces  the  copper 
oxide  to  copper  as  in  Exercise  n.  Then  disconnect  the  hydro- 
gen generator,  withdraw  the  heat,  and  allow  the  boat  to  remain 


Boat  filled  with  copper  oxide 


FIG.  45.  Apparatus  for  determining  the  formula  of  copper  oxide 

in  the  tube  until  cold.  Carefully  remove  the  boat  and  contents, 
and  weigh  to  o.oi  g.,  recording  the  weights  in  the  table. 
(a)  From  your  results  calculate  the  percentage  and  composition 
of  copper  oxide;  also  its  simplest  formula  as  directed  in  the 
following  table: 

Weight  of  boat    .    .    .    ,    .    .    .    . 

Weight  of  boat  plus  copper    oxide 

Weight  of  copper  oxide  taken  (calculate) 

Weight  of  boat  plus  residue  (copper)  left  after  heating 

Weight  of  copper  left  in  boat  (calculate) 

Percentage  of  copper  in  copper  oxide  (calculate  from 
formula) : 

Weight  of  copper  left  in  boat 

TTT  .  ,  . — 7 —          — . ,    ^  , —  X  100 

Weight  of  copper  oxide  taken 

Percentage  of  oxygen  in  copper  oxide   (ioo%  —  %  of 

copper  in  copper  oxide) 

Simplest  formula  (calculate  from  percentage  composition) 

[40] 


EXERCISE  19 

SOME  FURTHER  MANIPULATIONS 

Apparatus.  File;  forceps;  medicine  dropper;  (soo-cc.  flask  and 
cork  to  fit). 

Materials.  25  cm.  glass  tubing;  (50 cm.  glass  tubing;  5  cm.  soft 
rubber  tubing). 

1.  Devices   for  transfer- 
ring liquids  drop  by  drop. 

An  ordinary  medicine  drop- 
per (obtainable  at  any  drug 
store)  is  useful  for  transfer- 
ring a  few  drops  of  a  liquid 
from  one  vessel  to  another 
(Fig.  46).  A  glass  tube 
made  as  described  below  is 
often  of  more  service. 

Select  a  piece  of  glass  tub- 
ing about  2  5  cm.  long.  Heat 
the  glass  tube,  about  5  cm. 
from  one  end,  in  the  Bunsen 
flame  until  the  walls  of  the  heated  portion  thicken  and  the  size 
of  the  bore  diminishes  (A,  Fig.  47).  The  tube  must  be  con- 
stantly rotated  to  prevent  the  softened  portion  from  sagging. 
A, 


FIG.  46.   Transferring  a  liquid,  drop  by 
drop,  by  the  use  of  a  medicine  dropper 


FIG.  47.  Making  a  pipette 

Now  quickly  remove  the  tube  from  the  flame  and,  holding  it 
in  a  vertical  position,  gently  pull  the  ends  apart  until  the  bore 
is  of  the  desired  size  (B,  Fig.  47).  The  tube  is  then  cut  at  B, 

[41] 


and  both  ends  of  the  longer  piece  are  rounded  by  heating  in 
the  flame  as  directed  in  Exercise  3  (Fig.  16). 

To  transfer  a  liquid  from  one  vessel  into  another,  dip  the 
small  end  of  the  tube  below  the  liquid  to  be  transferred  and, 
placing  the  other  end  of  the  tube  in  the  mouth,  gently  suck  the 
liquid  up  in  the  tube  until  partially  filled.  Great  care  must  be 
taken  to  keep  the  end  of  the  tube 
well  below  the  surface  of  the 
liquid  while  it  is  being  drawn 
up  into  the  tube.  Moreover,  the 
tube  should  never  be  filled  more 
than  about  two  thirds;  other- 
wise the  liquid  may  be  drawn 
into  the  mouth  and  serious  re- 
sults follow..  When  the  tube  is 
partially  filled,  remove  the  end 
from  your  mouth  and  quickly 
press  your  forefinger  firmly  on 
the  upper  end  of  the  tube.  The 
liquid  in  the  tube  may  now  be 
transferred  to  another  vessel  by 
bringing  the  small  end  of  the 
tube  over  the  mouth  of  the  ves- 
sel and  then  slightly  decreasing 
the  pressure  of  the  finger  on  the  tube,  as  shown  in  Fig.  48. 
Such  a  tube  as  the  one  described  above  is  known  as  a  pipette. 

A  pipette  like  the  above,  but  graduated  in  cubic  centime- 
ters, serves  the  same  purpose  but  has  an  additional  advantage 
in  that  it  enables  one  to  tell  just  how  much  of  the  liquid  is 
transferred,  in  case  it  is  so  desired.  In  using  such  a  pipette,  the 
liquid  is  drawn  up  a  little  above  the  zero  point.  The  finger  is 
then  quickly  placed  against  the  end  of  the  pipette,  and  the 
liquid  allowed  to  flow  out  until  the  level  of  the  liquid  is 
just  at  the  zero  mark,  when  the  flow  is  stopped  by  increasing 
the  pressure  of  the  finger  against  the  end  of  the  pipette.  The' 

[421 


FIG.  48.  Transferring  a  liquid,  drop 
by  drop,  by  means  of  a  pipette 


liquid  may  then  be  allowed  to  drop  (or  flow)  into  any  vessel 
until  the  desired  amount  is  transferred.  The  flow  is  stopped  by 
increasing  the  pressure  of  the  finger  on  the  end  of  the  pipette. 


FIG.  49.  A  wash-bottle 


FIG.  50.  Method  of  using  a  wash-bottle 


2.  While  irot  essential,  the  student  will  find  a  so-called  wash-bottle 
very  convenient  for  adding  water  to  a  test  tube  or  beaker.  A  5oo-cc. 
flask  or  bottle  is  a  convenient  size.  The  apparatus  is  shown  in  Fig.  49. 
A,  B,  and  D  are  glass  tubes;  the  tube  D  is  drawn  out  to  a  small  bore 
(about  as  large  as  the  lead  in  a  lead  pencil).  The  tubes  B  and  D  are 
joined  by  the  rubber  tube  C.  The  bottle  is  filled  with  water,  and  a 
stream  is  projected  through  D  by  blowing  in  the  tube  A.  Fig.  50  shows 
the  method  of  using  the  bottle. 


[43 


EXERCISE  20 

THE  FORMATION  OF  CHARCOAL  AND  COKE 

Apparatus.  Hard-glass  test  tube  A  connected  as  shown  in  Fig.  51 ; 
B  is  a  test  tube  fitted  with  a  glass  tube  C  drawn  to  a  jet ;  large  beaker 
or  bottle. 

Materials.  Small  pieces  of  hard  wood  (sawdust  will  do)  sufficient  to 
half  fill  the  hard-glass  test  tube;  pieces  of  soft  coal  sufficient  to  half 
fill  the  test  tube. 

1.  Half  fill  the  tube  A  (Fig.  51)  with  hard-wood  splints  or 
sawdust  and  connect  it  as  shown  in  the  figure.  The  tube  B  is 
kept  cool  by  ice  water  in  the 
beaker  (or  bottle).  Heat  the 
wood  in  A,  gradually  increasing 
the  heat  until  no  further  change 
takes  place,  (a)  During  the 
heating,  test  the  gas  escaping 
from  the  jet  C  to  determine 
whether  it  will  burn. 

When  the  tube  A  is  cool,  re- 
move the  residue  and  note  its 
properties,  (b)  Is  it  combusti- 
ble? (c)  This  experiment  illus- 
trates the  manufacture  of  what 
substance  ? 

(d)  Note  the  odor  and  ap- 
pearance of  the  liquid  condensed 
in  B.  (e)  Name  two  important 
compounds  prepared  commer- 


FIG.  51.  Apparatus  for  heating  coal 
and  wood  in  the  absence  of  air  and 
for  collecting  the  products  formed 


daily  from  the  liquid  obtained  by  heating  hard  wood  in  the 
absence  of  air  (consult  text).  (/)  What  name  is  applied  to  the 
process  which  takes  place  when  any  combustible  substance, 
such  as  wood  or  coal,  is  heated  in  the  absence  of  air  ? 

[441 


2.  Repeat  the  experiment,  substituting  small  pieces  of  soft 
coal  for  the  wood,  (g)  Describe  the  results,  (h)  What  is  left 
in  the  tube  A  ?  The  black  liquid  condensed  in  B  is  known  as 
coal  tar  (consult  text),  (i)  Note  its  appearance  and  odor. 

EXERCISE  21 

A  FURTHER  STUDY  OF  CARBON 

Apparatus.  Hard-glass  test  tube;  porcelain  dish;  250-00.  flask; 
funnel. 

Materials.  3  g.  sugar ;  test  tube  one  fourth  full  of  boneblack ;  i  cc. 
litmus  solution  (R.  S.) ;  filter  paper ;  2  pieces  of  charcoal  the  size  of 
a  pea;  sulfuric  acid;  hydrochloric  acid. 

1.  Heat  2  or  3  g.  of  sugar  in  a  test  tube  until  no  further 
change  takes  place,    (a)  Note  the  results,    (b)  What  is  the 
residue  ? 

2.  Bring  a  cold  porcelain  dish  into  a  small  luminous  Bunsen 
flame,    (c)  Note  the  deposit,    (d)  What  is  this  form  of  carbon 
called?    (e)  In  what  other  forms  does  carbon  exist? 

3.  Fill  a  test  tube  about  one  fourth  full  of  boneblack  and 
transfer  it  to  a  small  flask ;  then  pour  over  the  boneblack  about 
50  cc.  of  water  to  which  has  been  added  a  few  drops  of  a  solu- 
tion of  litmus  or  indigo.   Thoroughly  mix  the  contents  of  the 
flask;  then  heat  gently  for  a  few  minutes  and  filter.    If  the 
filtrate  is  not  decolorized,  repeat  the  process,  using  more  bone- 
black.    (/)  What   is  the  composition  of  boneblack?    (g)  By 
what  other  name  is  it  known?    (h)  What  use  does  this  experi- 
ment suggest  for  it? 

4.  Place  a  piece  of  charcoal  in  a  test  tube  and  cover  it  with 
sulfuric  acid.    Does  any  action  seem  to  take  place?    Repeat, 
using  hydrochloric  acid.    (*')  Is  carbon  an  active  element  at 
ordinary  temperatures?    (;)  How  does  the  charring  of  wood 
preserve  it? 


[45 


EXERCISE  22 

THE  FORMATION  OF  CARBON  DIOXIDE  FROM  CARBON 
AND  OXYGEN 

Apparatus.  Apparatus  as  shown  in  Fig.  52.  Tube  A  and  bottle  B  are 
the  same  as  described  in  Fig.  43,  Exercise  16.  £  is  a  glass  tube,  and  C 
is  a  small  bottle  containing  about  10  cc.  of  limewater;  wing-top  burner. 

Materials.  15  cc.  limewater;  5  or  6  bits  of  charcoal  the  size  of  a  bean. 

1.  Test  for  carbon  dioxide.    Repeat  the  first  part  of  experi- 
ment 2,  Exercise  8.    Limewater  is  a  solution  of  ordinary  slaked 
lime  in  water.    Slaked  lime  is  a  white  solid  compound  of  the 
formula  CaO2H2  and  is  known  in  chemistry  as  calcium  hy- 
droxide.  The  change  in  the  limewater  produced  by  carbon 
dioxide  is  due  to  the  fact  that  the  carbon  dioxide  acts  upon 
the  calcium  hydroxide  in  the  solution  to  form  calcium  carbon- 
ate (CaCO3)  thus: 

CO2  +  CaO2H2  — >•  CaCO3  +  H2O 

The  calcium  carbonate  so  formed  is  a  white  solid  and  is  in- 
soluble in  water,  and  hence  separates,  as  fast  as  formed,  thus 
causing  the  limewater  to  become  cloudy.  This  reaction  serves 
as  a  good  test  for  carbon  dioxide. 

2.  Arrange  an  apparatus  as  shown  in  Fig.  52.    Place  in  the 
glass  tube  A  5  or  6  small  pieces  of  charcoal  and  connect  the 
apparatus  as  shown.    Pour  about  10  cc.  of  clear  limewater  into 
the  bottle  C.   Now  turn  the  water  cock  slightly  so  that  a  slow 
current  of  water  flows  into  the  bottle  B.   This  forces  the  air  in 
the  bottle  out  and  over  the  charcoal  in  A  and  finally  up  through 
the  limewater  in  C.    Continue  for  one  or  two  minutes,    (a)  Is 
there  any  change  in  the  appearance  of  the  limewater?    Turn 
off  the  water  and  gradually  heat  the  charcoal  in  tube  A  with 
the  wing-top  burner.   When  the  charcoal  is  hot  continue  the 
heating  and  again  turn  on  the  water  so  that  a  slow  current  of 

[46] 


air  is  forced  over  the  hot  charcoal,  (b)  Note  again  any  change 
in  the  limewater.  (c)  Is  carbon  dioxide  formed  in  the  experi- 
ment? (d)  What  is  its  source?  (e)  Write  the  equation  for  its 


Water  pipe 


Charcoal 


Air 


Limewater  — 


FIG.  52.  The  formation  of  carbon  dioxide  by  the  direct  combination  of  carbon 

and  oxygen 

formation.  (The  percentage  of  carbon  dioxide  in  the  air  is  so 
small  that  it  has  no  appreciable  effect  on  the  limewater  during 
the  time  of  the  experiment.) 

EXERCISE  23 
A  STUDY  OF  CARBON  DIOXIDE 

Apparatus.  Hydrogen  generator  with  connections,  as  shown  in 
Fig.  53  ;  three  25o-cc.  bottles ;  small  beaker  or  test  tube ;  hard-giass 
test  tube  ;  glass  rod. 

Materials.  5  pieces  of  marble  (size  of  walnuts);  hydrochloric  acid; 
splints;  limewater  (R.S.);  3  g.  of  copper  oxide  and  an  equal  bulk  of 
powdered  charcoal. 

1.  Usual  laboratory  method  for  preparing  carbon  dioxide. 
Place  some  pieces  of  marble  in  your  hydrogen  generator  and 

[47] 


connect  as  shown  in  Fig.  53. 
Add  water  through  the  funnel 
tube  until  the  marble  is  covered 
with  the  liquid ;  then  add  hydro- 
chloric acid,  a  few  drops  at  a 
time.  Carbon  dioxide  is  formed, 
passes  out  through  the  exit  tube, 
and  collects  in  the  bottle  B. 
Collect  three  bottles  of  the  gas 
by  displacement  of  air.  To  test 
when  filled,  hold  a  burning  splint 
at  the  mouth  of  the  bottle ;  the 
gas  will  extinguish  the  flame. 
(a)  Why  not  collect  the  gas  over 
water  as  in  the  case  of  oxygen 
and  hydrogen? 

2.  Properties   of   carbon   di- 
oxide,   (b)  Note  the  color  and 
odor  of  the  gas.    (c)  Pour  the 
gas  in  one  of  the  bottles  into 
another  bottle  filled  with  air, 
just  as  you  would  pour  water; 
then  test  for  the  gas  in  each  of 
the  bottles  with  a  lighted  splint. 

(d)  What  do  the  results  show  in 
reference  to  the  density  of  the 
gas  as  compared  with  air? 

3.  Chemical  conduct  of  carbon 
dioxide.     Introduce    a    lighted 
splint  into  a  bottle  of  the  gas. 

(e)  Is  the  gas  combustible?   (/) 
Is  it  a  supporter  of  combustion  ? 


o 


M 


7T 


FIG.  53.   The  preparation  of  carbon 

dioxide  by  the  action  of  hydrochloric 

acid  on  marble 


FIG.  54.  Testing  for  carbon  dioxide, 
using  a  drop  of  limewater 


Pour  5  cc.  of  clear  limewater  into  one  of  the  bottles;  then 
close  the  bottle  with  your  hand  and  shake  the  liquid,  (g)  Ex- 
plain the  results. 

[48] 


(h)  Mix  together  in  a  mortar  2  or  3  g.  of  black  copper  oxide 
and  an  equal  bulk  of  powdered  charcoal.  Transfer  to  a  hard- 
glass  test  tube  and  heat  gently.  The  copper  oxide  is  reduced 
to  copper,  the  oxygen  combining  with  the  hot  carbon  to  form 
carbon  dioxide.  Prove  the  presence  of  carbon  dioxide  in  the 
tube  as  follows :  Dip  the  end  of  a  glass  rod  into  limewater  and 
withdraw  it  gently  so  that  a  drop  of  the  clear  liquid  remains 
clinging  to  the  end  of  the  rod.  Lower  this  carefully  into  the  tube 
(Fig.  54).  Any  carbon  dioxide  present  will  cause  the  drop  of 
liquid  to  become  cloudy,  (i)  What  kind  of  an  agent  is  carbon? 
(;)  Briefly  summarize  the  properties  and  chemical  conduct  of 
carbon  dioxide. 

EXERCISE  24 

THE  PREPARATION  AND  PROPERTIES  OF  CARBON 
MONOXIDE  (OPTIONAL) 

Apparatus.  The  apparatus  shown  in  Fig.  55  :  A  is  a  funnel,  con- 
nected by  a  rubber  tube  with  a  piece  of  glass  tubing  which  passes 
through  a  stopper  into  the  25o-cc.  flask  B\  C  is  a  small  clamp;  pneu- 
matic trough ;  3  wide-mouthed  bottles ;  3  glass  plates  10  cm.  square. 

Materials.  Sulfuric  acid;  25  cc.  formic  acid  (50  per  cent);  lime- 
water  (R.  S.). 

Precaution.  Carbon  monoxide  is  a  nearly  odorless  and  very 
poisonous  gas  and  must  not  be  breathed.  All  the  experiments 
must  be  performed  in  the  hood.  After  the  gas  is  generated, 
pour  the  contents  of  the  generator  flask  into  a  sink  or  jar  in 
the  hood.  If  a  good  hood  is  not  available,  the  exercise  should 
be  omitted. 

1.  Preparation  of  carbon  monoxide.  Remove  the  stopper 
from  the  flask  B  (Fig.  55),  pour  in  15  cc.  of  sulfuric  acid,  and 
connect  the  apparatus  as  shown  in  the  figure.  Close  the  clamp 
C  and  partially  fill  the  funnel  A  with  the  formic  acid.  Now 
open  the  clamp  carefully  so  that  the  formic  acid  will  enter  the 
flask  drop  by  drop.  Allow  8  or  10  drops  to  flow  in ;  then  close 
the  clamp.  If  the  reaction  does  not  begin  (as  indicated  by 

[49] 


-Formic  acid 


FIG.  55.  Preparing  carbon  monoxide  by  the 
action  of  sulfuric  acid  on  formic  acid 


absence  of  effervescence  of  the  liquid  in  the  flask  and  escape  of 
gas  through  the  exit  tube),  heat  the  flask  very  gently  until  the 
reaction  starts;  then  open  the  clamp  again  and  admit  the 
formic  acid,  a  drop  at  a  time,  so  as  to  secure  a  regular  flow  of 
gas  from  the  flask.  If 
necessary,  add  more  for- 
mic acid  to  the  funnel  so 
as  to  keep  it  partially 
filled.  Collect  three  bot- 
tles of  the  gas  as  shown 
in  the  figure,  being  care- 
ful to  note  which  of  the 
bottles  was  the  first  one 
filled.  Close  the  clamp  so 
as  to  stop  further  gen- 
eration of  gas.  Slip  the 
glass  plates  over  the 
mouths  of  the  bottles  and 
remove  the  bottles  from  the  trough,  keeping  them  closed  with 
the  glass  plates,  (a)  Write  the  equation  for  the  preparation  of 
the  oxide  from  formic  acid. 

2.  Properties  and  chemical  conduct  of  carbon  monoxide. 
(b)  Has  the  gas  any  color?  (c)  Is  it  soluble  in  water? 

(d)  In  the  first  bottle  filled  under  i  is  the  gas  pure  carbon 
monoxide?  Test  it  by  introducing  a  burning  splint  into  the 
bottle ;  then  repeat  with  the  second  bottle,  (e)  Is  the  gas  com- 
bustible? (/)  Is  it  a  supporter  of  combustion?  Slip  the  glass 
plate  from  the  third  bottle  just  far  enough  to  pour  into  the 
bottle  5  cc.  of  clear  limewater ;  then  quickly  replace  the  glass 
plate  and,  holding  it  firmly  against  the  mouth  of  the  bottle, 
shake  the  contents  of  the  bottle.  Note  any  change  in  the  ap- 
pearance of  the  limewater.  Now  tip  the  bottle  as  far  as  possible 
without  spilling  the  limewater;  remove  the  glass  plate  and 
quickly  ignite  the  gas,  holding  the  bottle  in  this  position  so  that 
at  least  a  portion  of  the  combustion  product  may  be  retained  in 

[50] 


the  bottle.  When  the  flame  dies  out,  at  once  cover  the  mouth 
of  the  bottle  with  the  glass  plate  and  shake  the  contents. 
(g)  Record  and  explain  the  results. 

3.  (h)  Summarize  the  properties  and  chemical  conduct  of 
carbon  monoxide. 

EXERCISE  25 

THE  DETERMINATION  OF  THE  RELATIVE  VOLUMES  OF 
OXYGEN  AND  NITROGEN  IN  THE  AIR  (OPTIONAL) 

Apparatus.  Apparatus  shown  in  Fig.  56.  This  is  the  same  as  shown 
in  Fig.  43,  Exercise  16,  except  that  bottle  B  holds  not  less  than  2  1. 
while  bottle  C  holds  about  il.;  pneumatic  trough;  wing-top  burner; 
cork  to  fit  bottle  C ;  graduated  cylinder. 

Materials.  Copper  gauze  spiral  used  in  Exercise  16.  If  the  spiral  is 
oxidized,  reduce  it  as  directed  in  3,  Exercise  16. 

1.  To  determine  the  relative  volume  of  oxygen  and  nitrogen 
in  air,  a  definite  volume  of  air  is  passed  slowly  over  hot  copper, 


FIG.  56.  Determining  the  relative  volumes  of  oxygen  and  nitrogen  in  the  air 

which  combines  with  the  oxygen.  The  remaining  nitrogen  is 
collected  and  its  volume  measured.  The  experiment  is  the  same 
as  in  Exercise  16  except  that  the  volume  of  the  air  used  and 
that  of  the  resulting  nitrogen  are  both  measured, 

[51] 


Arrange  the  apparatus  as  shown  in  Fig.  56.  The  glass  tube 
A  contains  the  copper  spiral.  This  spiral  must  be  of  such  a 
size  that  it  will  just  enter  the  tube  and  yet  can  be  easily  pushed 
to  the  middle  part  of  the  tube  in  the  position  shown  in  the 
figure.  See  that  the  corks  and  tubes  are  air-tight.  At  the  be- 
ginning of  the  experiment  bottle  B  is  empty,  while  bottle  C  is 
completely  filled  with  water  and  inverted  in  G.  The  end  of 
tube  D  must  touch  the  bottom  of  bottle  B. 

When  the  apparatus  is  all  connected,  heat  the  copper  spiral 
in  tube  A  with  a  wing-top  burner  until  it  is  a  dull-red  heat,  and 
proceed  to  prepare  nitrogen  exactly  as  in  i,  Exercise  16.  Care 
must  be  taken  not  to  hasten  the  experiment,  as  some  of  the 
oxygen  may  not  combine  with  the  copper.  When  the  bottle 
C  is  nearly  filled,  turn  off  the  water,  withdraw  the  heat,  and 
disconnect  the  apparatus. 

Now  raise  (or  lower)  the  bottle  C  until  the  level  of  the  water 
inside  and  outside  the  bottle  is  the  same  (it  may  be  necessary 
to  add  more  water  to  the  trough) ;  then,  while  holding  it  in  this 
position,  push  the  cork  tightly  into  the  mouth  of  the  bottle. 
The  bottle  is  then  quickly  raised  from  the  trough  and  placed 
in  an  upright  position  on  your  desk. 

By  means  of  a  graduated  cylinder  measure  the  volume  of  the 
water  in  B ;  this  equals  the  volume  of  air  analyzed.  Record 
the  volume  in  the  table  below.  In  the  same  way  measure  the 
volume  of  the  nitrogen  in  C  (this  contains  about  i  per  cent  of 
rare  gases,  chiefly  argon)  and  record  it  in  the  table.  Then 
from  your  results  calculate  the  volumes  of  nitrogen  and  oxygen 
in  100  volumes  of  air,  recalling  that  100  volumes  of  air  contains 
approximately  i  volume  of  the  rare  gases  which  remain  mixed 
with  the  nitrogen. 

Volume  of  air  analyzed  (volume  of  water  in  B}  .  cc. 

Volume  of  nitrogen  and  rare  gases  in  the  volume  of 

air  analyzed  (volume  of  gas  in  C) cc. 

Volume  of  oxygen  in  air  analyzed   (difference  be- 
tween the  two  volumes  above) cc. 

[521 


Volume  of  nitrogen  and  rare  gases  in  100  volumes 

of  air  (calculate) cc. 

Volume  of  nitrogen  alone  in  100  volumes  of  air 
(equals  volume  of  nitrogen  and  rare  gases  in  100 
volumes  of  air  less  i  volume) cc. 

Volume  of  oxygen  in  100  volumes  of  air  (calculate)  cc. 

AVERAGE 
OF  RESULTS 
RESULTS       OBTAINED 
4  ACTUAL    OBTAINED       BY  CLASS 

Volumes  of  oxygen  in  100  volumes  of  air     21 
Volumes  of  nitrogen  in  100  volumes  of  air     78 

2.  (a)  What  are  the  sources  of  error  in  the  above  experi- 
ment? (b)  Why  is  no  attention  paid  to  the  temperature  of  the 
gases  or  to  the  barometric  pressure  ? 

EXERCISE  26 

A  STUDY  OF  SOLUTIONS 

Apparatus.  Test  tubes  ;  graduated  tube  or  pipette ;  funnel  and  filter 
paper;  loo-cc.  beaker;  watch  glass;  apparatus  shown  in  Fig.  57 
(A  is  a  250-cc.  flask,  B  is  a  thermometer,  and  C  is  an  open  glass  tube). 

Materials.  2  crystals  of  potassium  permanganate ;  0.2  g.  powdered 
calcium  sulfate ;  3  g.  common  salt ;  3  g.  potassium  nitrate  ;  i  g.  sugar ; 
10  cc.  carbon  tetrachloride  (R.  S.)  ;  i  cc.  oil  (cottonseed  or  olive)  ; 
log.  sodium  thiosulfate ;  sufficient  cotton  to  stopper  a  test  tube. 

1.  Nearly  fill  two  test  tubes  with  water  and  set  them  in  a 
rack.    Drop  into  each  a  small  crystal  of  potassium  permanga- 
nate.   Shake  the  contents  of  one  tube  and  repeat  the  shaking 
after  a  few  minutes,  but  do  not  move  the  other  tube,    (a)  At 
the  close  of  the  laboratory  period  note  the  appearance  of  the 
liquid  in  each  tube,    (b)  What  do  the  results  show? 

2.  Place  exactly  3  g.  of  common  salt  in  one  test  tube  and  an 
equal  weight  of  potassium  nitrate  in  another.   Add  to  each 
exactly  i  cc.  of  water  and  heat  each  tube  until  the  water  boils. 
If  the  solid  does  not  dissolve,  add  an  additional  cubic  centi- 
meter of  water  and  again  heat.    Repeat  until  the  solid  in  each 

[53] 


tube  is  dissolved,  avoiding  any  excess  of  water,  (c)  Compare  the 
approximate  solubilities  of  the  two  solids  in  the  boiling  water. 
(d)  Cool  the  solutions  in  each  tube  and  note  the  approxi- 
mate amounts  of  solids  separating,  (e)  Compare  the  results 
with  the  table  of  solubility  of  solids  given  in  the  Appendix 
of  text. 

3.  Introduce  about  0.5  g.  of  sugar  into  each  of  two  test  tubes. 
To  the  one  tube  add  5  cc.  of  water  and  to  the  other  5  cc.  of 
carbon  tetrachloride.    (/)   Shake  the  tubes 

gently  and  note  the  results.  Repeat  the  ex- 
periment, substituting  3  or  4  drops  of  an 
oil  for  the  sugar,  (g)  How  could  you  remove 
from  cloth  a  stain  due  to  a  sugar  sirup 
(molasses)  ?  to  oils  or  grease?  to  a  mixture 
of  the  two? 

4.  Introduce  into  a  test  tube  10  g.  of  so- 
dium thiosulfate  (ordinary  "hypo"  of  the 
photographer)  and  add  2  cc.  of  water.  Heat 
the  tube  gently  until  a  uniform  solution  is 
obtained,  care  being  taken  that  no  particles 
of  the  solid  remain  on  the  side  of  the  tube ; 
stopper  the  tube  loosely  with  a  plug  of  cot- 
ton and  set  it  aside  until  the  solution  is  cold.    FIG.  57.  Determining 
If  sufficient  care  has  been  taken,  no  solid    ^a^^£°J 
will  have  separated.  Now  remove  the  cotton 

plug  and  drop  into  the  solution  a  bit  of  the  solid  "hypo"  as 
large  as  a  pin  point,  (h)  Note  the  results  (hold  the  tube  in  a 
good  light)  and  explain. 

5.  By  means  of  the  apparatus  shown  in  Fig.  57,  determine 
the  boiling  point  of  a  saturated  solution  of  common  salt  in 
water,  keeping  the  bulb  of  the  thermometer  immersed  in  the 
solution,    (i)  Compare  the  boiling  point  of  the  solution  with 
that  of  pure  water. 


[54] 


EXERCISE  27 


S" 


If 


THE  DETERMINATION  OF  THE  SOLUBILITY  OF 
COMMON  SALT 

Apparatus.  6o-cc.  bottle ;  evaporating-dish  with  watch-glass  cover ; 
beaker,  tripod,  and  burner  as  shown  in  Fig.  58  ;  funnel  and  filter  paper ; 
thermometer. 

Materials.    15  g.  common  salt. 

1.  Introduce  about  15  g.  of  finely  powdered  common  salt  into 
a  6o-cc.  bottle  and  add  40  cc.  of  water.  Shake  the  mixture 
vigorously  and  set  aside  for  about 
ten  minutes,  repeating  the  shaking 
several  times  at  intervals  of  from 
one  to  two  minutes  so  as  to  form 
a  saturated  solution.  Take  the  tem- 
perature of  the  solution  and  record 
it  in  the  table  below. 

Accurately  weigh  to  o.oi  g.  a 
small  evaporating-dish  and  watch- 
glass  cover  and  record  the  weights 
in  the  table;  then  filter  into  the 
dish  about  20  cc.  of  the  saturated 
solution  of  salt  and  reweigh,  record- 
ing the  weights  in  the  table.  Remove 
the  watch  glass  and  place  the  dish  in  a  beaker  partially  filled 
with  water  as  shown  in  Fig  58.  Keep  the  water  in  the  beaker 
boiling.  After  the  solution  in  the  dish  has  evaporated  to  dry- 
ness,  remove  the  dish  and  place  it  on  a  ring  stand  (Fig.  24). 
Now  cover  the  dish  with  the  watch  glass  and  heat  it  directly 
with  the  burner,  regulating  the  flame  so  that  the  tip  barely 
touches  the  dish.  Continue  the  heating  until  all  the  moisture 
has  been  expelled  and  the  under  part  of  the  watch  glass  is  free 
from  moisture,  (a)  Why  use  the  watch-glass  cover? 

[55] 


FIG.  58.  Determining  the  solu- 
bility of  common  salt 


Now  withdraw  the  burner,  and  after  the  dish  is  cool  (room 
temperature)  weigh  the  dish,  residue,  and  cover,  recording  the 
weights  in  the  table. 

From  the  results  so  recorded,  fill  in  the  remaining  "blanks  in 
the  table  and  make  the  calculations  called  for.  (b)  Compare 
your  results  with  those  given  in  the  Appendix  of  the  text. 

Temperature  of  the  salt  solution / 

Weight  of  empty  dish  and  glass  cover g. 

Weight  of  dish,  salt  solution,  and  glass  cover   ....  g. 

Weight  of  salt  solution  taken g. 

Weight  of  dish,  residue  (salt)  left  after  evaporation  of 

water,  and  glass  cover g. 

Weight  of  water  present  in  solution  (loss  on  evaporation)  g. 

Weight  of  salt  dissolved  in  the  water  (residue)    .....  g. 
Weight  of  salt  that  will  dissolve  in  100  g.  of  water  at  the 

temperature  of  the  salt  solution  (calculate)     ....  g. 

EXERCISE  28 
THE  PREPARATION  AND  PROPERTIES  OF  CHLORINE 

Apparatus.  Test  tube  ;  apparatus  as  shown  in  Fig.  59  (A  is  a  250-00. 
flask  and  B  and  C  are  25o-cc.  bottles;  B  contains  some  sulfuric  acid); 
three  additional  25o-cc.  bottles  (dry)  ;  glass  plates. 

Materials.  25  g.  manganese  dioxide ;  hydrochloric  acid ;  sulfuric  acid ; 
bit  of  powdered  antimony ;  strips  of  colored  calico ;  piece  of  printed 
paper  (printer's  ink) ;  paper  written  over  with  ordinary  ink. 

Precaution.  All  the  following  experiments  must  be  performed 
in  the  hood,  and  great  care  must  be  taken  not  to  inhale  the 
chlorine. 

1.  Usual  laboratory  method  of  preparation.  (Two  students 
may  work  together.)  Arrange  an  apparatus  according  to 
Fig.  59.  Put  into  the  flask  from  20  to  2$g.  of  manganese 
dioxide.  Insert  the  cork  and  pour  150  cc.  of  hydrochloric  acid 
through  the  funnel  tube.  Shake  the  flask  so  as  to  mix  the  con- 
tents thoroughly.  Warm  gently,  applying  just  enough  heat  to 
cause  a  gentle  evolution  of  the  gas,  but  not  sufficient  to  boil 

[56] 


the  liquid.  Chlorine  is  set  free  and,  escaping  through  F,  bubbles 
through  the  sulfuric  acid  in  B  (which  removes  all  moisture) 
and  is  collected  in  C.  Fill  three  bottles  with  the  gas  (tell  when 
filled  by  the  yellowish  color  of  the  gas,  which  is  best  seen 
by  placing  a  white  cardboard  back  of  the  bottle).  Cover  the 
bottles  with  glass  plates 
and  set  them  aside. 

Last  of  all,  prepare 
some  chlorine  water  by 
bringing  the  exit  tube 
into  a  bottle  contain- 
ing 50  cc.  of  water  so 
that  the  gas  bubbles 
up  through  the  liquid. 
Continue  the  gentle 
heating  until  no  more 
chlorine  is  absorbed; 
then  cork  the  bottle, 
label  it,  and  set  it  aside 
in  a  dark  place  ((a) 


c 

H™"-       j 

Mr    n 

) 

n 

7 

I 

B 

c 

FIG.  59.  The  preparation  of  chlorine  by  the  ac- 
tion of  hydrochloric  acid  on  manganese  dioxide 


why  dark?)  for  future  use.     (b)  Write  the  equation  for  the 
reaction  that  takes  place  in  the  preparation  of  chlorine. 

2.  Chemical  conduct  of  chlorine.    Grind  a  fragment  of  anti- 
mony to  a  fine  powder  and  sprinkle  a  pinch  of  the  powder  into 
one  of  the  bottles  of  the  gas.    SbCl3  is  formed,    (c)  Note  the 
results  and  write  the  equation  for  the  reaction. 

Suspend  strips  of  colored  calico  in  a  bottle  of  the  gas ;  also 
two  strips  of  paper,  the  one  with  writing  in  ink  on  it,  the  other 
with  printing  (printer's  ink)  on  it.  Dip  similar  strips  of  calico 
and  paper  into  water  until  wet ;  then  suspend  these  in  another 
bottle  of  the  gas. 

(d)  Record  your  results  in  each  case,  (e)  What  part  does 
the  water  play  in  the  bleaching  ? 

3.  (/)  Summarize  the  properties  and  chemical  conduct  of 
chlorine. 

[57] 


EXERCISE  29 


THE  PREPARATION  AND  PROPERTIES  OF  HYDROGEN 
CHLORIDE  AND  OF  HYDROCHLORIC  ACID 

Apparatus.  Flask  and  bottle  connected  as  shown  in  Fig.  60  (this 
is  same  as  shown  in  Fig.  59,  except  that  the  bottle  B  contains  water  and 
the  glass  tube  extending  into  the  bottle  B  does  not  touch  the  water  in  B) : 
two  250-cc.  bottles  (dry) ;  large  beaker ;  medicine  dropper. 

Materials.  Dilute  sulfuric  acid  prepared  (care)  by  slowly  pouring 
(i  or  2  cc.  at  a  time  with  constant  stirring)  30  cc.  of  the  concentrated 
acid  into  10  cc.  water  ;  50  g.  sodium  chloride  ;  splint ;  blue  litmus  paper. 

1.  Usual  laboratory  method  of  preparing  hydrogen  chloride. 
(Hood.)  Arrange  the  apparatus  as  shown  in  Fig.  60.  The 
bottle  B  contains  water.  Care 
must  be  taken  that  the  glass 
tube  in  the  bottle  B  does  not 
quite  touch  the  surface  of  the 
water  in  the  bottle.  Put  about 
50  g.  of  common  salt  into  the 
flask  A,  insert  the  cork,  pour 
the  cold  dilute  sulfuric  acid 
through  the  funnel  tube,  and 
mix  the  contents  by  a  gentle 
motion  of  the  flask;  then  con- 
nect the  flask  with  F  as  shown 
in  the  figure.  After  two  or  three 
minutes  warm  gently  with  a 
small  flame.  Notice  the  currents 
in  the  water  in  B.  (a)  What 
causes  them  ?  When  the  gas  is  evolved  regularly,  disconnect  the 
generator  flask  at  D  long  enough  to  collect  two  bottles  (dry)  of 
the  gas  by  displacement  of  air  as  in  Fig.  53  (tell  when  filled  by 
the  fog  formed  at  mouth  of  bottle).  Cover  these  tightly  with 
dry  glass  plates  and  set  them  aside ;  then  connect  the  generator 

[58] 


FIG.  60.   The  preparation  of  hydro- 
gen chloride  by  the  action  of  sul- 
furic acid  on  sodium  chloride 


with  B  again  and  continue  to  apply  a  gentle  heat  as  long  as 
any  gas  is  evolved,    (b)  Write  the  equation  for  the  reaction. 

2.  Properties  and  chemical  conduct  of  hydrogen  chloride. 
(c)  What  is  the  color  of  the  gas  (examine  that  in  flask  A}? 
Test  the  gas  in  one  of  the  bottles  with  a  lighted  splint,    (d)  Is 
it  combustible?    (e)  Is  it  a  supporter  of  combustion?    Fill  a 
large  beaker  with  water.   Now  uncover  the  remaining  bottle, 
invert  it,  and  at  once  bring  its  mouth  under  the  surface  of  the 
water  in  the  beaker  and  hold  it  in  this  position  for  two  or  three 
minutes.    (/)  Describe  the  results,    (g)  What  does  the  experi- 
ment prove?  (h)  Why  not  extend  the  tube  in  bottle  B  (Fig.  60) 
to  the  bottom  of  the  bottle  ? 

3.  Properties  of  hydrochloric  acid.    Put  a  drop  of  the  solu- 
tion from  bottle  B  on  a  bit  of  blue  litmus  paper,    (i)  Note  the 
result.    (;)  Pour  2  drops  of  the  solution  into  5  cc.  of  water  and 
taste  a  drop,    (k)  Perform  a  test-tube  experiment  to  prove  the 
presence  of  chlorine  in  the  acid ;  (/)  also  one  to  prove  the  pres- 
ence of  hydrogen,    (m)  How  does  the  solution  compare  with 
the  hydrochloric  acid  on  your  desk? 

4.  ( n )  State  clearly  the  difference  between  the  terms  hydrogen 
chloride  and  hydrochloric  acid. 

EXERCISE  30 

SODIUM ;  SODIUM  HYDROXIDE 

Apparatus.  Evaporating-dish  with  glass-plate  cover ;  forceps ;  glass 
rod. 

Materials.    Bit  of  sodium  as  large  as  a  pea ;  red  litmus  paper. 

1.  Properties  of  sodium.  Recall  experiment  i,  Exercise  9. 
Obtain  from  your  instructor  a  bit  of  sodium.  Hold  the  sodium 
with  the  forceps  on  a  glass  plate,  cut  it,  and  note  the  rapidity 
with  which  the  freshly  cut  surface  is  tarnished.  Half  fill  your 
evaporating-dish  with  water;  then  drop  into  this  a  bit  of 
sodium  as  large  as  a  pea,  quickly  cover  the  dish  with  the  glass 
plate,  and  leave  it  covered  until  the  sodium  has  all  disappeared. 

[591 


(a)  Is  sodium  heavier  or  lighter  than  water?  (b)  Write  the 
equation  for  the  reaction  between  sodium  and  water,  (c)  What 
is  the  composition  of  the  liquid  in  the  evaporating-dish  ? 

2.  Properties  of  sodium  hydroxide.  Place  a  drop  of  the 
liquid  in  the  evaporating-dish  on  a  piece  of  red  litmus  paper. 

(d)  Contrast  its  action  with  that  of  hydrochloric  acid  on  litmus 
paper  (Exercise  29).    Mix  i  or  2  drops  of  the  solution  with 
5  cc.   of  water  and  taste  a  drop  of  the  resulting  solution. 

(e)  Contrast  with  the  taste  of  hydrochloric  acid  (Exercise  29). 
Evaporate  the  solution  to  dryness.    (/)  What  is  the  residue? 

EXERCISE  31 

THE  PROPERTIES  OF  ACIDS,  BASES,  AND  SALTS 

Apparatus.  Small  beaker ;  stirring-rod ;  evaporating-dish  ;  ring  stand ; 
test  tubes ;  medicine  dropper. 

Materials.  A  few  drops  of  each  of  the  following  acids:  hydrochloric, 
sulfuric,  nitric,  acetic  ;  solutions  of  the  following  bases  :  sodium  hydrox- 
ide, potassium  hydroxide  (R.  S.),  calcium  hydroxide  (R.  S.);  strips  of 
blue  and  of  red  litmus  paper ;  hydrochloric  acid. 

1.  Acids.    Recall  the  properties  of  hydrochloric  acid  (Exer- 
cise 29).    Prepare  a  dilute  solution  of  each  of  the  following 
acids  by  adding  2  or  3  drops  of  the  acid  to  about  10  cc.  of  water 
in  a  test  tube  and  mixing  thoroughly:  hydrochloric,  sulfuric, 
nitric,  acetic. 

By  means  of  a  clean  glass  rod  transfer  a  drop  of  each  of  the 
dilute  solutions  to  a  piece  of  blue  litmus  paper,  (a)  Note  the 
result  in  each  case,  (b)  Taste  one  drop  of  the  dilute  solutions 
(rinse  the  mouth  with  water  after  tasting). 

Compare  the  formulas  of  the  acids,  (c)  In  what  respect  are 
the  acids  similar  in  composition  ? 

2.  Bases.    In  a  similar  way  try  the  effect  on  red  litmus  paper 
of  a  solution  of  each  of  the  following  bases :  sodium  hydroxide 
(recall  Exercise  30),  potassium  hydroxide,  calcium  hydroxide. 
(d)  Do  they  affect  the  blue  litmus  paper?    (e)  Taste  a  drop 
of  the  calcium  hydroxide  solution. 

[60] 


Compare  the  formulas  of  the  bases.  (/)  In  what  respect  are 
the  bases  similar  in  composition? 

3.  Salts.    Dilute  5  cc.  of  the  ordinary  laboratory  solution  of 
sodium  hydroxide  (i  part  of  the  hydroxide  by  weight  to  10 
parts  of  water)  with  an  equal  volume  of  water.    To  this  solu- 
tion add  4  or  5  drops  of  the  ordinary  concentrated  hydrochloric 
acid.    Stir  the  resulting  solution  with  a  glass  rod  and  test  its 
action  on  blue  and  on  red  litmus  paper,    (g)  Has  it  acid  or 
basic  properties? 

Now  continue  to  add  the  acid  drop  by  drop  until  the  result- 
ing solution  is  neutral  (that  is,  has  no  effect  on  either  blue  or 
red  litmus  paper)  or  is,  at  most,  slightly  acid,  (h)  Write  the 
equation  for  the  reaction  between  the  hydroxide  and  the  acid. 
Now  pour  the  solution  into  an  evaporating-dish  and  evaporate 
to  dryness.  (i)  What  compound  remains?  (;)  Taste  it. 

(k)  What  is  the  name  given  to  the  compounds  formed  by 
the  action  of  acids  on  bases  ? 

4.  (/)  Characterize  acids  and  bases  as  to  composition;  as  to 
their  action  on  litmus ;  as  to  taste ;  as  to  their  interaction  with 
each  other. 

EXERCISE  32 

THE  RATIO  OF  ACID  TO  BASE  IN  NEUTRALIZATION 

Apparatus.  2  burettes  and  supports,  as  shown  in  Fig.  61;  2  small 
beakers  and  stirring-rod;  graduated  pipette. 

Materials.  Sodium  hydroxide  solution  prepared  by  adding  20  cc.  of 
the  laboratory  reagent  to  100  cc.  water;  i  cc.  sulfuric  acid  (use  gradu- 
ated pipette)  added  to  100  cc.  water  and  mixed  thoroughly;  a  few  drops 
of  a  phenolphthalein  solution  (R.  S.). 

(Two  students  may  work  together.) 

1.  Rinse  out  a  burette,  first  with  distilled  water  and  then 
with  a  little  of  the  solution  of  sodium  hydroxide.  Support  the 
burette  (Fig.  61)  and  pour  into  it  the  hydroxide  solution  until 
the  level  of  the  liquid  is  i  or  2  cm.  above  the  zero  mark.  Turn 
the  stopcock  and  let  the  solution  slowly  flow  out  until  the  bot- 

[61] 


torn  of  the" curved  surface  (Fig.  3)  of  the  liquid  in  the  burette 
is  on  a  level  with  the  zero  mark.  In  a  similar  way  fill  a  second 
burette  with  the  acid  solution. 

Now  let  exactly  1 5  cc.  of  the  acid 
solution  flow  into  a  small  beaker,  add 
two  drops  of  phenolphthalein  solution, 
and  run  in  2  or  3  cc.  of  the  hydroxide 
solution.  Notice  that  where  the  liquids 
come  in  contact  a  reddish  color  is  pro- 
duced, which  disappears  quickly  on 
stirring.  Run  in  more  of  the  solution, 
a  little  at  a  time,  until  the  color  fades 
slowly,  and  then  a  drop  at  a  time  until 
the  entire  liquid,  on  stirring,  remains 
colored  faintly  pink.  This  marks  ap- 
proximately the  point  of  neutralization. 
Record  the  number  of  cubic  centimeters 
of  the  hydroxide  solution  used  in  the 
table  below.  Repeat  the  experiment, 


FIG.  61.   Graduated  tubes 

(burettes)  for  measuring 

volumes  of  liquids 


using  30  cc.  of  the  acid,  and  again  record  in  the  table  the 
number  of  cubic  centimeters  of  the  hydroxide  solution  used 
to  neutralize  the  acid. 

First  Experiment 

Volume  of  hydroxide  solution  used  to  neutralize  15  cc. 

of  the  acid cc. 

Volume  of  hydroxide  solution  required  to  neutralize 

i  cc.  of  acid  (calculated) 


Second  Experiment 

Volume  of  hydroxide  solution  used  to  neutralize  30  cc. 

of  the  acid 

Volume  of  hydroxide  solution  required  to  neutralize 

i  cc.  of  acid  (calculated) 

2.  (a)  What  do  the  results  of  the  experiments  prove  ? 


cc. 


cc. 


cc. 


62] 


EXERCISE  33 

CARBONIC  ACID  AND  ITS  SALTS  (CARBONATES) 

Apparatus.  Hydrogen  generator,  as  used  for  preparing  carbon 
dioxide  in  Exercise  23  ;  small  beaker ;  5  test  tubes. 

Materials.  Pieces  of  marble  for  generating  carbon  dioxide;  hydro- 
chloric acid ;  blue  litmus  paper ;  5  cc.  sodium  hydroxide  solution  diluted 
with  10  cc.  water ;  i  g.  of  each  of  the  common  carbonates,  such  as 
sodium  carbonate,  magnesium  carbonate,  calcium  carbonate ;  25  cc. 
limewater  (R.  S.). 

1.  When  carbon  dioxide  is  passed  into  water,  a  portion  of  the 
gas  combines  with  the  water  to  form  a  compound  known  as  car- 
bonic acid  ( H2O  +  CO2  - — >•  H2CO3 ) .    This  is  a  very  weak  acid 
and  so  unstable  that  it  can  be  obtained  only  in  dilute  solutions. 
If  an  attempt  is  made  to  concentrate  the  solution  by  evaporat- 
ing the  water,  the  acid  decomposes  again  into  water  and  carbon 
dioxide.   The  salts  of  this  acid  (the  carbonates),  on  the  other 
hand,  are  stable,  and  many  of  them  are  important  and  common 
compounds.   Thus,  calcium  carbonate   (CaCO3)   is  the  chief 
constituent  of  limestone,  while  sodium  carbonate  (Na2CO3) 
is  common  washing  soda. 

Generate  carbon  dioxide  (Fig.  53)  and  pass  the  gas  through 
25  cc.  of  water,  (a)  Taste  the  liquid,  (b)  Is  the  acid  formed 
strong  enough  to  affect  blue  litmus  paper? 

2.  Pass  carbon  dioxide  through  5  cc.  of  a  solution  of  sodium 
hydroxide  until  the  gas  is  no  longer  absorbed   (the  carbon 
dioxide  combines  with  the  water  present  to  form  carbonic  acid. 
This  acid  then  reacts  with  the  base,  sodium  hydroxide,  to 
form  a  salt  and  water).    Evaporate  the  solution  to  dryness. 
(c)  What  is  the  product  left  in  the  dish?    (d)  Write  the  equa- 
tion for  its  formation,    (e)  Could  a  solution  of  sodium  hy- 
droxide be  used  in  place  of  a  solution  of  calcium  hydroxide 
(limewater)  in  testing  for  carbon  dioxide? 

[63] 


3.  Examine  the  physical  properties  of  such  carbonates  as  are 
available.  (/)  What  ones  are  soluble  in  water  (p.  304  of  text)  ? 
Test  the  action  of  hydrochloric  acid  or  sulfuric  acid  on  each  by 
adding  i  or  2  drops  of  the  acid  to  o.i  g.  of  the  carbonate  in  a 
test  tube.  What  evidences  have  you  that  a  gas  is  evolved? 
(g)  Test  the  gas  to  determine  whether  or  not  it  is  carbon 
dioxide  (Fig.  54).  All  carbonates  evolve  carbon  dioxide  when 
treated  with  hydrochloric  or  sulfuric  acid.  This  reaction  serves 
as  a  good  test  for  carbonates. 

EXERCISE  34 

A   METHOD   FOR  DETERMINING  WHETHER  A  GIVEN 
LIQUID  IS  A  CONDUCTOR  OF  ELECTRICITY  (OPTIONAL) 

Apparatus.  Current  from  electric-lighting  system ;  apparatus  as 
shown  in  Fig.  62.  B  is  a  plug  for  connecting  the  apparatus  with  any 
ordinary  electric-lighting  system.  C  is  an  ordinary  incandescent  lamp. 
A  is  a  small  bottle  (60  cc.).  The  wires  are  the  common  insulated  cop- 
per wires,  but  the  insulation  is  removed  from  that  portion  of  the  wires 
which  extend  inside  the  bottle  A.  The  apparatus  can  be  purchased 
from  supply  houses,  but  is  easily  made. 

Materials.  5  g.  common  salt ;  5  g.  sugar ;  5  cc.  sodium  hydroxide 
solution  added  to  20  cc.  water ;  3  cc.  sulfuric  acid  added  to  20  cc. 
water ;  tap  or  well  water ;  10  cc.  hydrochloric  acid ;  (10  cc.  of  a  benzene 
solution  of  hydrogen  chloride). 

1.  Obtain  from  your  instruc- 
tor the  apparatus  shown  in 
Fig.  62.  Polish  the  ends  (elec- 
trodes) of  the  copper  wires 
that  extend  into  the  bottle  A 

with    emery   paper    until    they     FIG.  62.    Apparatus   for  determining 

are  bright  and  free  from  oxide.    whether  or  not  a  lic*uid  is  a  conduc- 

.......  ,          ,  tor  of  electricity 

At  the  beginning  of  each  ex- 
periment see  that  the  electrodes  are  bright  and  dry  and  that 
the  bottle  is  also  perfectly  clean  and  dry.  Unscrew  a  lamp  C 
from  a  convenient  socket  in  the  laboratory,  screw  it  loosely 

[641 


into  the  socket  on  your  apparatus,  and  attach  the  apparatus  to 
the  empty  socket  on  the  lighting  system  by  means  of  the  ex- 
tension cord  and  plug  B.  Every  time  a  change  is  to  be  made 
in  the  cell,  loosen  the  lamp  C  in  the  socket,  and  do  not  screw 
it  down  to  make  contact  until  all  the  connections  of  the  cell 
have  been  arranged. 

2.  Partly  fill  the  bottle  A  with  dry,  powdered  salt,  dip  the 
electrodes  into  the  powder,  arrange  the  connections  at  the 
binding-posts,  and  screw  down  the  lamp  C.    (a)  Does  the  salt 
conduct  the  electric  current? 

3.  (b)  Repeat  2,  using  distilled  water. 

4.  (c)  Repeat  2,  using  a  solution  of  the  salt  in  distilled  water. 

5.  (d)  Test  the  conductivity  of  the  following  substances  and 
interpret  the  results :  dry,  powdered  sugar ;  a  solution  of  sugar 
in  water;  tap  or  well  water;  distilled  water  containing  a  few 
drops  of  sulfuric  acid;  distilled  water  containing  a  few  drops 
of  hydrochloric  acid ;  a  solution  of  sodium  hydroxide,  (e)  How 
do  we  account  for  the  fact  that  solutions  of  some  compounds 
conduct  the  electric  current  while  others  do  not? 

6.  Secure  from  your  instructor  10  cc.  of  a  benzene  solution  (benzene 
is  inflammable)  of  hydrogen  chloride  (easily  prepared  just  as  an  aque- 
ous solution  is  prepared,  as  described  in  Exercise  29).  (/)  Test  its  con- 
ductivity, (g)  Test  its  effect  on  blue  and  on  red  litmus  paper,  (h}  Does 
it  dissolve  zinc  ?    (i)  How  do  you  account  for  the  difference  in  properties 
between  a  benzene  solution  of  hydrogen  chloride  and  an  aqueous  solu- 
tion of  the  same  gas  ? 

7.  (;)  Define  the  terms  acid,  base,  and  salt  from  the  stand- 
point of  the  ionization  theory. 


65] 


EXERCISE  35 


THE  DISPLACEMENT  OF  METALS  FROM  THEIR 
COMPOUNDS  (THE  DISPLACEMENT  SERIES) 

Apparatus.    4  test  tubes;  test-tube  rack. 

Materials.  4  clean  and  bright  strips  each  of  zinc  and  copper 
(i  cm.  x  10  cm.);  0.5  g.  lead  nitrate  dissolved  in  10  cc.  water;  0.5  g. 
copper  nitrate  dissolved  in  10  cc.  water;  0.5  g.  mercuric  nitrate  dis- 
solved in  10  cc.  water;  3  cc.  sulfuric  acid  dissolved  in  10  cc.  water. 

1.  Pour  into  separate  test  tubes  to  a  depth  of  4  or  5  cm.  solu- 
tions of  the  following  compounds :  lead  nitrate,  dilute  sulfuric 
acid,  copper  nitrate,  mer- 
curic nitrate.  Set  the 
tubes  in  a  rack  in  the 
order  given  above  and 
label  them  A,  B,  C,  and 
D  respectively  (Fig.  63). 

Now  place  in  each  tube 
a  strip  of  zinc.  (It  is  con- 
venient to  have  a  strong 
thread  attached  to  the 
upper  part  of  each  strip 
so  that  the  strip  may 
easily  be  withdrawn  from 
the  tube.  The  strips 
should  be  only  partly  immersed  in  the  solution.)  Note  any 
change  taking  place  in  the  appearance  of  the  zinc. 

After  twenty  minutes  withdraw  the  strips  and  wipe  them 
on  a  piece  of  white  paper,  (a)  Note  any  evidence  tending  to 
show  that  the  zinc  has  displaced  the  lead,  hydrogen,  copper, 
and  mercury  from  their  compounds.  (Metals  in  a  very  finely 
divided  form  are  black,  as  a  rule.)  (b)  Account  for  any  change 
in  the  color  of  the  copper  nitrate  solution. 

[661 


FIG.  63.  Testing  the  action  of  different  metals 
on  solutions  of  salts 


2.  Repeat  experiment  i,  substituting  for  the  zinc  a  strip  of 
copper,  (c)  Contrast  the  results  obtained  with  those  obtained 
in  i.  (d)  How  do  you  account  for  the  change  in  the  color  of  the 
solution  of  mercuric  nitrate  after  the  addition  of  the  copper 
strip?  (e)  Are  your  results  in  accord  with  the  table  of  the 
displacement  series  given  on  page  161  of  text? 

EXERCISE  36 

THE  PREPARATION  AND  PROPERTIES  OF  AMMONIA 

Apparatus.  Test  tube;  hard-glass  test  tube  fitted  with  cork  and 
tubing  as  shown  in  Fig.  64  ;  graduated  tube  ;  three  250-00.  wide-mouthed 
bottles  ;  2  pieces  of  window  glass  ;  large  beaker  ;  glass  rod. 

Materials.  Solution  of  sodium  hydroxide  (the  ordinary  desk  solution 
will  serve)  ;  9  g.  ammonium  chloride  ;  15  g.  powdered  calcium  hydroxide 
(slaked  or  hydrated  lime,Ca(OH)2);  strip  of  red  litmus  paper;  strip 
of  blue  litmus  paper  ;  hydrochloric  acid  ;  splint. 

1.  Dissolve  0.5  g.  of  ammonium  chloride  in  3  or  4  cc.  of  water 
in  a  test  tube  and  heat  to  boiling,  (a)  Note  the  odor. 

Now  add  3  cc.  of  a  solution  of  sodium  hydroxide  to  the  hot 
solution  of  ammonium  chloride  and  continue  the  heating. 
(b)  Again  note  the  odor,  (c)  Moisten  a  strip  of  red  litmus 
paper  and  hold  it  at  the  mouth  of  the  tube  but  not  in  contact 
with  it. 

(d)  Dip  the  end  of  a  glass  rod  in  a  concentrated  solution  of 
hydrochloric  acid  and  hold  it  in  the  mouth  of  the  test  tube. 
Dense  white  fumes  of  ammonium  chloride  (NH4C1)  are  formed. 

(e)  Complete  the  following  equations: 

NH4Cl  +  NaOH  -  >• 
NH3  +  H2O  —  >- 


NH4OH 

2.  Usual  laboratory  method  for  preparing  ammonia.  This 
differs  from  the  method  used  in  i  only  in  the  fact  that  the 
less  expensive  calcium  hydroxide  (slaked  lime)  is  substituted 

[67] 


for  the  sodium  hydroxide.  The  form  of  apparatus  used  is  shown 
in  Fig.  64.  The  bottle  B  (250-00.)  contains  2500.  of  water. 
The  glass  tube  C  extends  through  a  hole  in  a  cardboard  resting 
on  the  mouth  of  the  bottle.  The  end  of  the  tube  must  just 
touch  the  water  in  the  bottle. 

Place  in  the  tube  A  a  mixture  of  15  g.  of  powdered  slaked 
lime  and  8  g.  of  ammonium  chloride.  Connect  the  tube  as 
shown  in  Fig.  64  and 
heat  the  mixture  gently, 
beginning  with  that 
portion  near  the  mouth 
of  the  tube  and  gradu- 
ally extending  the  heat 
to  the  other  portions. 
As  soon  as  the  gas 
is  evolved  freely  (as 
shown  by  the  bubbles 
at  the  end  of  the  tube 
C),  bring  the  tube  C  to 
an  upright  position,  as 


FIG.  64.  Preparing  ammonia  by  heating  a  mix- 
ture of  ammonium  chloride  and  calcium  hy- 
droxide (slaked  lime) 


shown  in  the  dotted 
lines,  and  collect  two 
bottles  of  the  gas.  To 
do  this,  bring  the  bottles  successively  down  over  the  exit  tube, 
leave  each  in  this  position  until  a  drop  of  hydrochloric  acid 
on  the  end  of  a  glass  rod  fumes  strongly  when  held  at  the 
mouth  of  the  bottle ;  then  withdraw  the  bottle,  cover  its  mouth 
with  a  glass  plate,  and  set  it  aside,  mouth  downward. 

When  both  bottles  are  filled,  bring  the  tube  C  into  the  bottle 
B  again  and  continue  to  heat  the  mixture  gently  as  long  as  any 
gas  is  generated.  (/)  Write  the  equations  for  all  the  reactions 
involved. 

3.  Properties  and  chemical  conduct  of  ammonia,  (g)  Note 
the  color  and  odor  of  the  gas.  (h)  Is  it  heavier  or  lighter 
than  air?  (i)  Test  a  bottle  of  the  gas  with  a  burning  splint. 

[68] 


Devise  a  simple  experiment  for  finding  out  whether  or  not 
ammonia  is  soluble  in  water.  After  your  method  is  approved 
by  the  instructor,  try  it  out  with  the  remaining  bottle  of  the 
gas.  (;)  Describe  the  results. 

4.  Properties  of  ammonium  hydroxide,  (k)  Note  the  odor  of 
the  liquid  in  the  bottle  B.    (I)  Try  its  effect  on  blue  and  on 
red  litmus  paper,    (m)  How  does  it  compare  with  the  aqua 
ammonia  of  the  druggist  in  its  odor  and  its  action  on  litmus? 
(n)  Does  the  gas  combine  with  the  water  or  is  it  simply  dis- 
solved in  the  water  ?    (0)  Give  reasons  for  your  answer.   Now 
neutralize  the  liquid  with  hydrochloric  acid  and  evaporate  just 
to  dry  ness  (Fig.  58).    (/>)  Compare  the  residue  with  the  am- 
monium chloride  used  in  experiment  i. 

5.  (q)  Summarize  the  properties  and  chemical  conduct  of 
ammonia. 

EXERCISE  37 
THE  PREPARATION  AND  PROPERTIES  OF  NITRIC  ACID 

Apparatus.  Glass  retort  (150-00.),  test  tube,  and  beaker  (500-00.), 
arranged  as  shown  in  Fig.  65;  funnel;  evaporating-dish. 

Materials.  12  g.  sodium  nitrate  ;  10  cc.  sulfuric  acid ;  small  piece  of 
tin;  small  strip  of  copper. 

Caution.  The  students  must  remember  that  both  sulfuric  acid 
and  nitric  acid  are  very  corrosive  and  must  exercise  care  in 
handling  them. 

1.  Preparation  of  nitric  acid.  Arrange  an  apparatus  like  that 
shown  in  Fig.  65.  Put  in  the  retort  A  about  12  g.  of  sodium 
nitrate  and  10  cc.  of  sulfuric  acid,  pouring  the  latter  through 
a  funnel  placed  in  the  tubulus  B  of  the  retort.  Heat  the  mix- 
ture gently  with  a  small  flame.  Nitric  acid  is  set  free,  distills 
over,  and  is  condensed  in  the  test  tube  C,  which  is  kept  cold 
by  being  partly  immersed  in  ice  water  in  the  beaker  D. 
(a)  Write  the  equation  for  the  reaction  between  the  sulfuric 
acid  and  the  sodium  nitrate. 

[69] 


2.  Properties  and  chemical  conduct  of  nitric  acid.  When 
nitric  acid  is  heated,  a  part  of  it  is  decomposed  into  water, 
nitrogen  dioxide,  and  oxygen.  On  this  account  it  is  a  good 
oxidizing  agent.  To  test  its  oxidizing  properties,  put  a  small 
piece  of  tin  in  a  test  tube,  cover  it  with  a  little  nitric  acid,  and 
gently  heat  (hood). 
The  white  product 
formed  is  composed 
mainly  of  tin  and  oxy- 
gen, the  latter  being 
supplied  by  the  nitric 
acid. 

Pure  nitric  acid  is 
colorless,  (b)  How  do 
you  account  for  the 
color  of  the  acid  which 


FIG.  65.   The  preparation  of  nitric  acid  by  the 
action  of  sulfuric  acid  on  sodium  nitrate 


you  have  prepared? 

Place  a  small  strip 
of  copper  in  an  evaporating-dish  (hood)  and  add  some  of  the 
acid  you  have  prepared,  a  few  drops  at  a  time,  until  the  copper 
is  just  dissolved.  Evaporate  the  solution  to  dryness  (Fig.  58). 
(c)  Note  the  appearance  of  the  residue,  (d)  Since  copper  is 
below  hydrogen  in  the  displacement  series,  how  do  you  account 
for  the  fact  that  nitric  acid  dissolves  the  metal  (consult  text)  ? 
Save  the  residue  in  the  dish  ((e)  what  is  it?)  for  use  in  the 
following  exercise. 


[70 


EXERCISE  38 

THE  PROPERTIES  OF  THE   SALTS  OF  NITRIC  ACID 
(NITRATES) 

Apparatus.  Evaporating-dish  containing  the  copper  nitrate  prepared 
in  Exercise  37 ;  5  test  tubes  in  test-tube  rack;  medicine  dropper. 

Materials.  Strip  of  copper ;  crystal  of  lead  nitrate ;  crystals  of  such 
nitrates  as  are  available,  including  sodium  nitrate  and  potassium  nitrate ; 
sulfuric  acid ;  2  g.  ferrous  sulfate  dissolved  in  10  cc.  water. 

1.  Properties  of  nitrates.   Heat  the  dish  (hood)  containing 
the  copper  nitrate,  prepared  in  Exercise  37,  with  a  small  flame. 
(a)  Note  the  color  of  the  gas  evolved,  also  the  color  of  the 
residue,    (b)  Compare  the  residue  with  the  crust  of  copper 
oxide  prepared  by  holding  a  strip  of  copper  in  the  tip  of  a  flame. 

Place  a  crystal  of  lead  nitrate  in  the  evaporating-dish  and 
heat  gently,  (c)  Compare  with  the  results  obtained  in  i. 

Place  a  small  crystal  of  such  nitrates  as  are  available  in  your 
laboratory  in  separate  test  tubes  and  test  their  solubility  in 
water;  (d)  What  nitrates  are  insoluble  in  water  (consult  text)  ? 

2.  How  to  detect  the  presence  of  a  nitrate.    Dissolve  a  crystal 
of  sodium  nitrate  in  2  or  3  cc.  of  water  in  a  test  tube,  add  drop 
by  drop  an  equal  volume  of  sulfuric  acid,  mixing  the  liquids 
as  the  acid  is  added,  then  cool  the  mixture.    The  sulfuric  acid 
acts  on  the  nitrate,  liberating  nitric  acid.   Now  tip  the  tube 
slightly  and  gently  pour  2  or  3  cc.  of  the  solution  of  ferrous 
sulfate  down  the  side  of  the  tube,  so  that  it  floats  on  the  heavier 
liquid,  and  set  the  tube  aside,  being  careful  not  to  mix  the  two 
liquids.   A  brown  ring  soon  forms  where  the  liquids  meet. 
Repeat  the  experiment,  using  potassium  nitrate.    This  is  a  good 
test  for  nitrates.  The  brown  ring  is  due  to  the  presence  of  a  com- 
pound formed  by  the  action  of  ferrous  sulfate  on  nitric  acid. 

Secure  one  or  more  compounds  from  your  instructor  and  test 
to  see  if  they  are  nitrates,  (e)  Record  your  results. 

[71] 


EXERCISE  39 

THE  PREPARATION  AND  PROPERTIES  OF  SOME  OF 
THE  OXIDES  OF  NITROGEN 

Apparatus.  Hard-glass  test  tube,  with  delivery  tube,  as  used  in  pre- 
paring oxygen  (Fig.  28);  3  wide-mouthed  bottles  (250-00.);  pneumatic 
trough;  hydrogen  generator  (Fig.  35)52  glass  plates. 

Materials.  8g.  ammonium  nitrate;  wooden  splints;  5  small  strips  of 
copper;  10  cc.  nitric  acid. 

1.  Nitrous  oxide.   Put  6  or  8  g.  of  ammonium  nitrate  in  the 
hard-glass  test  tube  used  in  the  preparation  of  oxygen  (Fig.  28). 
Attach  a  delivery  tube  and  heat  gently,  applying  no  more  heat 
than  is  necessary  to  cause  a  slow  evolution  of  the  gas. 

As  soon  as  the  gas  is  regularly  evolved,  collect  two  or  three 
bottles  of  it  over  water.  Notice  the  water  deposited  on  the 
sides  of  the  test  tube,  (a)  What  is  the  source  of  it?  (b)  Note 
the  color,  odor,  and  taste  of  the  gas.  (c)  Test  it  with  a  glowing 
splint,  (d)  Account  for  the  result,  (e)  How  can  you  distin- 
guish between  nitrous  oxide  and  oxygen  ? 

2.  Nitric  oxide  and  nitrogen  dioxide.    Put  a  few  pieces  of 
copper  in  your  hydrogen  generator  (hood)  (Fig.  35),  just  cover 
them  with  water,  and  add  2  or  3  cc.  of  nitric  acid.    Collect 
over  water  two  bottles  of  the  evolved  gas,  adding  more  nitric 
acid  to  the  liquid  in  the  generator  if  necessary  to  secure  the 
required  amount  of  gas. 

(/)  Compare  the  color  of  the  gas  in  the  generator  with, 
that  collected  in  the  bottles  and  account  for  any  difference. 
(g)  Write  the  equations  for  all  the  reactions  involved. 

(h)  Uncover  one  of  the  bottles  of  the  gas  and  account  for 
the  result,  (i)  Test  the  gas  in  the  second  bottle  with  a  burning 
splint.  (;)  Which  is  the  more  stable,  nitrous  oxide  or  nitric 
oxide?  (k)  Give  reasons  for  your  answer. 


72] 


EXERCISE  40 
THE  PROPERTIES  AND  CHEMICAL  CONDUCT  OF  SULFUR 

Apparatus.  3  test  tubes;  smallest-sized  beaker;  magnifying-glass ; 
porcelain  crucible ;  pipestem  triangle ;  large  beaker ;  steel  forceps. 

Materials.  5  cc.  carbon  disulfide ;  20  g.  powdered  brimstone  ;  strip 
of  copper  foil  (0.5  x  3  cm.) ;  5  g.  iron  powder. 

1.  Properties  of  sulfur,  (a)  Examine  the  physical  prop- 
erties of  a  piece  of  brimstone.  Pour  2  or  3  cc.  of  carbon  disul- 
fide (hood)  over  2  g.  of  powdered  brimstone  in  a  test  tube. 
(Keep  carbon  disulfide  away  from  flame  and  do  not  inhale 
the  vapor.}  Cover  the  mouth  of  the  tube  with  the  thumb  and 
shake  the  contents  gently  until  the  sulfur  is  dissolved,  adding 
more  carbon  disulfide  if  necessary.  Pour  the  clear  solution 
into  a  small  beaker,  cover  the  beaker  loosely  with  a  filter  paper, 
and  set  it  aside  in  the  hood.  The  carbon  disulfide  soon  evapo- 
rates, the  sulfur  being  deposited  in  crystals,  (b)  Examine  these 
with  a  magnifying-glass,  noting  their  general  shape. 

Half  fill  a  test  tube  with  powdered  brimstone  and  heat  it 
gently  until  the  sulfur  is  just  melted,  (c)  Note  the  properties 
of.  the  liquid.  Now  apply  a  stronger  heat  and  observe  that  the 
liquid  becomes  darker,  and  at  a  certain  temperature  (200°- 
250°)  is  so  thick  that  the  tube  may  be  inverted  without  spill- 
ing it.  Finally,  increase  the  heat  until  the  sulfur  boils  (448°), 
and  then  slowly  pour  the  boiling  liquid  into  a  beaker  of  cold 
water,  (d)  Examine  the  product,  (e)  What  name  is  given  to 
this  form  of  sulfur  ?  Expose  it  to  the  air  for  an  hour  or  longer. 
(/)  Have  its  properties  remained  unchanged? 

Fill  a  porcelain  crucible  with  powdered  brimstone  and  apply 
a  very  gentle  heat  until  the  sulfur  is  just  melted.  Withdraw 
the  flame  and  examine  the  liquid  carefully  as  it  cools.  Crystals 
soon  begin  to  form  on  the  surface  of  the  melted  sulfur,  rapidly 
extending  from  the  circumference  toward  the  center.  Before 

[73] 


they  reach  the  center  grasp  the  crucible  with  the  forceps  and 
quickly  pour  off  the  remaining  liquid  and  examine  the  crystals. 
(g)  Compare  them  in  shape  with  those  deposited  from  the 
solution  of  carbon  disulfide. 
(h)  In  how  many  forms  have 
you  obtained  sulfur? 

2.  Chemical  conduct  of  sul- 
fur. Burn  a  small  piece  of 
sulfur.  (/')  Notice  the  appear- 
ance of  the  burning  sulfur 
and  the  odor  of  the  gas  (SO2) 
formed. 

Boil  a  little  sulfur  in  a  test 
tube  and  drop  a  small  strip 
of  hot  copper  foil  (Fig.  66) 
in  the  boiling  liquid.  (;)  Is 
there  any  visible  evidence  of 


FIG.  66.  Dropping  a  piece  of  copper  foil 
into  boiling  sulfur 


a  chemical  change?  (k)  What 
is  formed  ? 

Grind  together  5  or  6  g.  of  iron  powder  with  an  equal  weight 
of  sulfur,  transfer  the  mixture  to  a  test  tube,  and  heat  it  strongly 
in  a  Bunsen  flame.  (/)  Describe  the  results.  Retain  the  solid 
formed  for  use  in  the  following  exercise. 

3.  (m)  Summarize  the  properties  and  conduct  of  sulfur. 

EXERCISE  41 

THE  PREPARATION   AND   PROPERTIES   OF  HYDROGEN 

SULFIDE 

Apparatus.  Hydrogen  generator  and  tubes,  as  shown  in  Fig.  67  ; 
2  wide-mouthed  bottles  (250-00.  and  6o-cc.)  ;  funnel ;  evaporating-dish. 

Materials.  10  g.  ferrous  sulfide  (in  pieces  as  big  as  a  bean  or  larger)  ; 
20  cc.  hydrochloric  acid  added  to  20  cc.  water ;  blue  and  red  litmus 
papers  ;  silver  coin ;  filter  paper. 

1.  (Hood.)  Attach  a  delivery  tube  to  the  hydrogen  gen- 
erator, as  shown  in  Fig.  67.  Put  into  the  generator  A  a  few 

[74] 


o 


JR 


TT 


pieces  of  ferrous  sulfide  (FeS)  and  insert  the  stopper.  Now 
pour  a  little  water  through  the  funnel  tube  of  the  generator 
until  the  end  of  the  tube  just  dips  below  the  surface  of  the 
water ;  then  pour  in  a  few  cc.  of  the  hydrochloric  acid,  adding 
more  from  time  to  time,  if  necessary,  to  maintain  a  gentle 
evolution  of  the  gas.  (a)  Write  the  equation.  The  gas  escapes 
into  the  bottle  B,  which  gradually  becomes  filled,  (b)  Very 
cautiously  note  the  odor  (Caution.  The  gas  is  poisonous  if 
inhaled  in  concentrated  form)  ; 
(c)  also  note  the  color  of  the 
evolved  gas.  Continue  the  evo- 
lution of  the  gas  until  it  is 
ignited  by  a  flame  held  at  the 
mouth  of  the  bottle  B.  (d)  Ac- 
count for  the  deposit  formed 
on  the  sides  of  the  bottle  when 
the  gas  burns. 

2.  Replace  the  bottle  B  with 
a  6o-cc.  bottle  half  filled  with 
water,  and  allow  the  gas  from 
the  generator   (add  more  acid 
if  necessary)  to  bubble  through 
the  water  for  one  or  two  min- 
utes.  Test  the  resulting  solution  with  blue  and  with  red  litmus 
paper,    (e)  What  is  the  solution  called?    (/)   How  does  it 
compare  with  the  so-called  "sulfur  water"  of  many  springs? 
(g)  Drop  a  silver  coin  into  the  solution  and  account  for  the 
results,   (h)  Why  do  certain  foods  blacken  silver  spoons? 

3.  Transfer  to  a  test  tube  a  piece  of  the  solid  formed  in  2, 
Exercise  40,  by  heating  a  mixture  of  iron  and  sulfur  and  add  i  or 
2  cc.  of  dilute  hydrochloric  acid,     (i)  What  gas  is  evolved 
(cautiously  note  the  odor)  ?    (j)  Account  for  its  formation. 

4.  (k)  Summarize  the  properties  of  hydrogen  sulfide. 


B 


FIG.  67.  Preparing  hydrogen  sulfide 

by  the  action  of  hydrochloric  acid 

on  ferrous  sulfide 


[75 


EXERCISE  42 


THE  PREPARATION  AND  PROPERTIES  OF  THE  SALTS  OF 
HYDROSULFURIC  ACID  (SULFIDES) 

Apparatus.  Hydrogen  generator  and  connections,  as  shown  in  Fig.  68  ; 
6  test  tubes ;  beaker ;  funnel ;  watch  glass  ;  iron  spoon. 

Materials.  Ferrous  sulfide  and  dilute  hydrochloric  acid,  as  used  in 
Exercise  41 ;  separate  solutions  of  silver  nitrate,  copper,  sulfate,  cad- 
mium chloride,  lead  nitrate,  and  sodium  chloride,  made  by  dissolving 
about  0.5  g.  of  the  solid  in  5  cc.  water  (solutions  on  reagent  shelf  may 
be  used)  ;  sulfuric  acid ;  lead  acetate  (R.  S.)  ;  5  g.  sulfur;  3  g.  lime. 

1.  Preparation  of  sulfides.  (Hood.)  Charge  the  hydrogen 
sulfide  generator  as  in  Exercise  41  and  pass  a  few  bubbles  of 
the  gas  (Fig.  68)  through  each  of  the  following  solutions: 
silver  nitrate,  copper  sulfate, 
cadmium  chloride,  sodium 
chloride,  lead  nitrate.  The 
exit  tube  C,  through  which 
the  gas  bubbles  into  the  solu- 
tions, must  be  thoroughly 
cleaned  each  time,  (a)  Note 
the  color  of  the  precipitate 
obtained  in  each  case.  Write 
the  equations  for  the  reac- 
tions involved,  (b)  Do  any 
of  the  solutions  fail  to  give  a 
precipitate?  (c)  How  do  you 


O 


A 


FIG.  68.  Preparing  sulfides  of  the  metals 

by  passing  hydrogen  sulfide  through  a 

solution  of  their  salts 


account  for  this  ? 

Intimately  mix  5  g.  of  sul- 
fur with  3  g.  of  powdered  lime.  Transfer  to  a  beaker  and  add 
i5occ.  of  water.  Stir  the  mixture  and  heat  just  to  boiling  for 
ten  minutes.  Now  fill  a  test  tube  with  the  resulting  mixture, 
cork  the  tube  loosely,  and  set  it  aside  until  the  beginning  of 
the  next  laboratory  period;  then  examine,  (d)  Describe  the 

[76] 


results.,    (e)  For  what  is  the  solution  used  (consult  text)  ? 
(/)  What  is  its  composition  (consult  text)  ? 

2.  Test  for  hydrogen  sulfide.    Dip  a  strip  of  filter  paper  into 
a  solution  of  lead  acetate.    Remove  the  cork  from  the  hydrogen 
sulfide  generator  and  insert  the  paper  for  a  moment,   (g)  Note 
and  account  for  the  results.    This  serves  as  a  convenient  test 
for  the  gas.   (h)  What  property  would  also  serve  to  detect  the 
gas  if  present  in  any  marked  quantity? 

3.  Test  for  sulfides.    Prepare  some  dilute  sulfuric  acid  by 
adding   (care)    2   or  3   drops  of  the  acid  on  your  desk  to 
i  cc.  of  water.    Place  a  small  bit  of  ferrous  sulfide   (FeS) 
on  your  watch  glass  and  moisten  it  with  the  dilute  acid. 
Cautiously  note  the  odor.   Most  of  the  sulfides  when  treated 
in  this  way  evolve  hydrogen  sulfide,  which  can  be  detected 
by  its  odor  and  by  its  action  on  paper  moistened  with  lead 
acetate  solution. 

All  sulfides  when  heated  in  air  evolve  sulfur  dioxide  (formed 
by  the  combustion  of  the  sulfur  present),  which  has  the  char- 
acteristic odor  of  burning  sulfur,  (i)  Heat  a  little  ferrous  sul- 
fide on  an  iron  spoon  in  the  flame  of  the  burner  and  note 
the  odor. 

EXERCISE  43 
SULFUR  DIOXIDE  AND  SULFUROUS  ACID 

Apparatus.  25o-cc.  flask  fitted  with  funnel  tube  and  glass  exit  tube, 
as  shown  in  Fig.  69 ;  3  bottles  (250-00.) ;  watch  glass  ;  medicine  dropper. 

Materials.  10  g.  copper ;  sulfuric  acid  ;  splint ;  blue  litmus  paper ; 
crystal  of  sodium  sulfite ;  strips  of  colored  calico  or  a  red  flower. 

1.  (a)  By  what  method  have  you  already  prepared  sulfur 
dioxide  ? 

2.  Preparation  of  sulfur  dioxide  and  sulfurous  acid  by  the 
action  of  sulfuric  acid  on  copper  (laboratory  method).    Place 
about  10  g.  of  copper  turnings  or  small  pieces  of  sheet  copper 
in  a  generator  arranged  as  in  Fig.  69.   Add  250:.  of  concen- 

[77] 


trated  sulfuric  acid  and  apply  a  gentle  heat.  As  soon,  as  the 
action  begins,  lower  the  flame,  regulating  it  so  as  to  obtain 
a  uniform  evolution  of  the  gas.  Collect  two  bottles  of  the  gas 
by  displacement  of  air;  then 
cause  it  to  bubble  through 
2  5  cc.  of  water  as  long  as  any 
is  dissolved,  (b)  Write  the 
equation  for  the  reaction  of 
sulfuric  acid  on  copper. 

3.  Properties  and  chemical 
conduct  of  sulfur  dioxide  and 
sulfurous  acid,  (c)  Note  the 
odor  of  the  gas.  (d)  Is  the  gas 
combustible?  (e)  Test  with 
blue  litmus  paper  the  liquid 
formed  by  passing  the  gas 


FIG.  69.    Preparing  sulfur  dioxide  by 
the  action  of  sulfuric  acid  on  copper 


through  water.    (/)   Does  the 

gas  combine  with  the  water  or 

simply  dissolve  in  it  ?  (g)  Give 

reasons  for  your  answer,   (h)  Immerse  in  the  liquid  some  small 

strips  of  colored  calico  or  some  petals  of  a  red  flower  and  note 

any  results. 

4.  Salts  of  sulfurous  acid;  the  sulfites.    Place  a  small  crystal 
of  some  sulfite,  such  as  sodium  sulfite  (Na2SO3),  on  a  watch 
glass  and  moisten  it  with  2  or  3  drops  of  sulfuric  acid,    (i)  Note 
the  odor  of  the  evolved  gas.    All  sulfites  evolve  sulfur  dioxide 
when  treated  with  sulfuric  acid.   This  reaction  serves  as  a  good 
test  for  sulfites.    It  also  serves  as  a  method  for  preparing 
sulfur  dioxide. 

5.  (j)  Enumerate  three  methods  for  preparing  sulfur  dioxide. 

6.  (k)  Briefly  summarize  the  properties  and  chemical  con- 
duct of  sulfur  dioxide. 


[78] 


EXERCISE  44 

A  STUDY  OF  SULFURIC  ACID 

Apparatus.    Stirring-rod ;   test  tubes ;  medicine  dropper. 

Materials.  Sulfuric  acid  ;  0.5  g.  sugar ;  wooden  splint ;  a  small  piece 
of  zinc;  piece  of  charcoal  the  size  of  a  pea;  barium  chloride  solution 
(R.S.);  hydrochloric  acid. 

1.  Properties  and  chemical  conduct  of  sulfuric  acid.    By 
means  of  a  glass  rod  place  i  or  2  drops  of  sulfuric  acid  on  a 
wooden  splint,  and  after  two  or  three  minutes  (a)  note  the 
results  produced.    Put  about  0.5  g.  of  sugar  in  a  test  tube  and 
add  3  or  4  drops  of  the  acid,    (b)  Note  and  account  for  the 
change  produced. 

Heat  a  bit  of  charcoal  in  a  test  tube  with  i  or  2  cc.  of  con- 
centrated sulfuric  acid,  (c)  What  gas  is  evolved  (odor)  ? 
(d)  Account  for  its  formation,  recalling  that  carbon  has  a 
strong  affinity  for  oxygen. 

In  the  preparation  of  hydrogen  (Exercise  9)  the  directions 
called  for  dilute  sulfuric  acid,  (e)  Determine  whether  the 
concentrated  acid  will  do  as  well,  by  pouring  2  or  3  cc.  of  the 
acid  over  a  bit  of  zinc  in  a  test  tube  (if  no  action  takes  place, 
heat  the  mixture  gently  and  test  the  gas  evolved  both  by  a 
lighted  splint  and  by  noting  its  odor).  (/)  Account  for  the 
difference  in  the  action  on  zinc  between  the  dilute  and  the 
concentrated  acid. 

Recall  the  action  of  sulfuric  acid  on  sodium  chloride  (Exer- 
cise 29)  and  on  sodium  nitrate  (Exercise  37).  (g)  What  prop- 
erty of  sulfuric  acid  enables  it  to  be  used  in  the  preparation 
of  other  acids  ? 

2.  Test  for  sulfuric  acid.    Add  3  drops  of  sulfuric  acid  (care) 
to  5  cc.  of  water  in  a  test  tube.   To  this  add  a  few  drops  of  a 
solution  of  barium  chloride.    The  chloride  reacts  with  the  acid 
to  form  barium  sulfate  (BaSO4),  which  is  insoluble  and  hence 

[79] 


separates,  as  fast  as  formed,  as  a  white  precipitate,  (h)  Write 
the  equation  for  the  formation  of  the  barium  sulfate.  Now  add 
to  the  precipitate  4  or  5  drops  of  hydrochloric  acid,  (i)  Does 
the  precipitate  dissolve  ?  The  formation  of  a  white  precipitate 
with  barium  chloride,  insoluble  in  hydrochloric  acid,  consti- 
tutes a  good  test  for  sulfuric  acid. 

EXERCISE  45 
SALTS  OF  SULFURIC  ACID  (SULFATES) 

Apparatus.    6  test  tubes. 

Materials.  Crystals  or  small  amounts  (o.i  g.)  of  the  sulfates  avail- 
able in  the  laboratory;  2  cc.  barium  chloride  solution  (R.S.);  hydro- 
chloric acid. 

1.  Examine  the  physical  properties  of  such  sulfates  as  are 
available.    Test  the  solubility  of  each  in  water  by  adding 
about  0.5  g.  of  the  sulfate  to  a  test  tube  nearly  filled  with  water 
and  shaking  the  mixture,     (a)   What  sulfates  are  insoluble 
(consult  text)  ? 

2.  Prepare  a  dilute  solution  of  different  sulfates  by  dissolv- 
ing a  crystal  of  each  in  2  or  3  cc.  of  water.   Add  to  each  i  drop 
of  barium  chloride  solution.    Now  add  2  drops  of  hydrochloric 
acid  to  the  mixture  in  each  tube,    (b)  Does  the  precipitate  dis- 
solve?  All  soluble  sulfates  give  in  solution  a  white  precipitate 
(BaSO4)  with  barium  chloride  solution,  which  precipitate  is 
insoluble  in  hydrochloric  acid.   This  reaction  serves  as  a  good 
test  for  sulfates. 

It  will  be  noted  that  both  sulfuric  acid  and  its  salts  give  with  barium 
chloride  the  same  product;  namely,  a  white  precipitate  of  barium  sul- 
fate. This  is  evident  from  the  following  facts :  It  will  be  recalled  that 
both  acids  and  salts  are  ionized  in  solution.  In  the  case  of  sulfuric  acid 
and  sulfates,  ions  are  formed  as  follows  : 

H2S04 >-H+  H+  +  S04— 

Na~S04 >-Na+,  Na+-f  SO, — 

CuS04 >•  Cu+  +  +  S04 — 

[801 


Likewise,  barium  chloride  solution  contains  the  ions  Ba++  and  Cl~, 
Cl~.  Now  when  a  solution  of  barium  chloride  is  mixed  with  any  so- 
lution containing  the  SO4 —  ion,  the  two  ions  Ba++  and  S04 —  unite 
to  form  the  insoluble  BaSO4,  which  precipitates;  hence  the  reaction 
proceeds  to  completion  (paragraph  2,  p.  179  of  text).  The  barium 
chloride  test  is  therefore  really  a  test  for  the  presence  of  the  SO4  ion. 
Since  only  sulfuric  acid  and  its  salts  give  this  ion,  however,  it  is  cus- 
tomary to  say  that  it  is  a  test  for  sulfuric  acid  and  the  sulfates. 

EXERCISE  46 

THE  PREPARATION  AND  PROPERTIES  OF  HYDROGEN 
FLUORIDE 

Apparatus.  Piece  of  window  glass;  small  lead  dish  (laboratory  outfit). 
Materials.    2  or  3  small  pieces  of  paraffin  (size  of  a  pea)  ;  3  g.  fluor- 
ite  (CaF2)  ;  sulfuric  acid. 

Precaution.  Hydrogen  fluoride  is  very  corrosive  and  must 
not  be  inhaled ;  neither  must  its  solution  be  brought  in  contact 
with  the  skin. 

1.  Place  some  pieces  of  paraffin  on  a  glass  plate  and  gently 
warm  the  plate  over  a  small  flame.  When  the  paraffin  is  melted, 
tilt  the  plate  about  so  as  to 
form  a  uniform  layer  of  the 
wax.  When  the  wax  is  cold, 
scratch  your  name  through 
the  wax  with  a  pin  (Fig. 
70).  Place  3  g.  of  fluorite 
in  a  lead  dish  and  add  suffi-  FlG.  70.  Preparing  a  glass  plate  for  etching 
cient  sulfuric  acid  to  make 

a  paste  of  it.  Cover  the  dish  tightly  with  the  waxed  side  of 
the  glass  plate  and  set  it  in  the  hood  for  an  hour ;  then  scrape 
off  the  paraffin  and  examine  the  glass,  (a)  Describe  the-results. 
(b)  Write  the  equations  for  all  the  reactions  involved. 


[81 


EXERCISE  47 


THE  PREPARATION  AND  PROPERTIES  OF  BROMINE 

Apparatus.  Retort,  test  tube,  and  beaker,  as  shown  in  Fig.  71  ;  funnel. 

Materials.  3  g.  sodium  bromide  or  potassium  bromide  ;  4  g.  manga- 
nese dioxide  ;  10  cc.  sulfuric  acid  dissolved  in  40  cc.  water  ;  strips  of  col- 
ored calico;  i  cc.  carbon  tetrachloride  ;  i  cc.  of  chlorine  water  (R.  S.). 

Precaution.     The  vapor  of  bromine  must  not  be  inhaled. 

1.  Preparation  of  bromine.  Put  into  the  retort  (A,  Fig.  71) 
a  mixture  of  2  g.  of  potassium  bromide  or  of  sodium  bromide 
and  4  g.  of  manganese  dioxide,  and  add  to  this  through  a  funnel 
the  cold  dilute  solution  of  sulfuric  acid  prepared  as  directed 
above.  Shake  the  re- 
tort so  as  to  mix  the 
contents  thoroughly. 
Note  that  the  test- 
tube  receiver  C  con- 
tains sufficient  water 
to  allow  the  end  of 
the  retort  to  dip  just 
below  its  surface. 

Now  heat  the  retort 

gently.     The  bromine 

is    liberated    and    dis- 
JM1  /    \    TIT  •, 

tills  over,    (a)  Write 


FIG.  71.    Preparing  bromine  by  the  action  of  Sul- 

*ur*c  ac^  on  a  Imxture  °f  manganese  dioxide 

and  sodium  bromide 


the  equation  for  the  reaction.  Continue  the  heating  until  all 
the  bromine  has  distilled  over.  Remove  the  stopper  from  the 
retort  before  the  heat  is  withdrawn. 

2.  Properties  of  bromine,  (b)  Note  the  properties  of  the 
bromine  collected  in  the  bottom  of  the  test  tube,  (c)  Has  any 
dissolved  in  the  water  (note  the  color  of  the  water)  ?  (d)  What 
property  is  implied  in  the  name  of  the  element?  Add  about 
0.5  cc.  of  the  aqueous  bromine  solution  to  i  cc.  of  carbon 

[82] 


tetrachloride.  Note  that  the  bromine  solution  does  not  mix  with 
the  carbon  tetrachloride  but  floats  on  top  of  it.  Now  place 
your  thumb  over  the  mouth  of  the  tube  and  shake  the  mixture 
vigorously  for  a  few  seconds ;  then  set  the  tube  aside  until  the 
two  liquids  separate.  The  bromine,  being  more  soluble  in  carbon 
tetrachloride  than  in  water,  is  taken  up  by  the  former  liquid. 
(e)  What  is  the  color  of  the  carbon  tetrachloride  solution? 

Test  the  bleaching  property  of  bromine  by  immersing  strips 
of  colored  cloth  in  the  aqueous  solution.  (/)  How  does  it 
compare  with  chlorine  as  a  bleaching  agent? 

3.  Action  of  chlorine  on  sodium  bromide.   Dissolve  a  bit  of 
sodium  bromide  as  large  as  a  grain  of  wheat  in  2  cc.  of  water. 
(g)  What  is  the  color  of  the  solution  ?   Now  add  to  the  solution 

1  cc.  of  chlorine  water.   The  chlorine  displaces  bromine  from 
sodium  bromide  just  as  zinc  displaces  copper  from  a  copper 
salt  (see  displacement  series,  Exercise  35).    (h)  Account  for 
the  change  in  the  color  of  the  solution. 

4.  (i)  Summarize  the  properties  and  conduct  of  bromine. 

EXERCISE  48 
THE  PREPARATION  AND  PROPERTIES  OF  IODINE 

Apparatus.  Beaker  (250-00.)  and  evaporating-dish  as  shown  in  Fig.  72 ; 

2  test  tubes  ;  stirring-rod. 

Materials.  4  g.  sodium  iodide  (or  potassium  iodide) ;  4  g.  powdered 
manganese  dioxide;  sulfuric  acid;  2  cc.  alcohol;  2  cc.  carbon  tetra- 
chloride; 10  cc.  chlorine  water  (R.S.);  10  cc.  starch  solution  (R.S.). 

1.  Preparation  of  iodine.  Intimately  mix  on  paper  3  g.  of 
powdered  sodium  iodide  and  2  g.  of  manganese  dioxide.  Place 
the  mixture  in  the  bottom  of  the  beaker  A  (Fig.  72).  Nearly 
fill  the  evaporating-dish  B  with  cold  water.  Now  remove  the 
evaporating-dish,  add  5  cc.  of  sulfuric  acid  to  the  mixture  in 
the  beaker,  and  stir  it  through  the  mixture ;  then  at  once  place 
the  evaporating-dish  back  on  the  beaker  as  shown  in  the  figure. 
Apply  a  very  gentle  heat  to  the  mixture  in  the  beaker,  moving 

[83] 


B 
A 


\NaI 
\MnO, 
\H2S04 


the  burner  about  so  as  to  heat  the  mixture  uniformly.  Note  the 
colored  vapor  formed,  which  slowly  condenses  on  the  sides  of 
the  beaker  and  the  bottom  of  the  evaporating-dish  in  the  form 
of  grayish-black  crystals.  Finally  withdraw  the  flame  and  let 
the  beaker  stand  until  the  vapor  is  nearly  all  condensed. 
(a)  What  is  the  substance?  (b)  Account  for  its  formation. 
(c)  What  property  does  the  name  of 
the  substance  suggest  ? 

2.  Properties  of  iodine.  By  means 
of  a  glass  rod  transfer  a  little  of  the 
iodine  collected  on  the  sides  of  the 
beaker  in  i  to  a  test  tube  half  filled 
with  water,  and  shake  it  for  one  or 
two  minutes,  (d)  Is  the  iodine  soluble 
in  water  (note  the  color  of  the  water) 
to  any  extent?    (Set  the  tube  and 
contents  aside  for  use  in  4,  below.) 
Repeat,   using   alcohol   in   place   of 
water,   (e)  Compare  the  solubility  of 
the  iodine  in  water  and  in  alcohol. 
(/)  What  is  the  solution  in  alcohol 

called.?   Dissolve  a  small  bit  of  the  iodine  in  2  cc.  of  carbon 
tetrachloride.     (g)  Note  the  color  of  the  solution. 

3.  Action  of  chlorine  on  sodium  iodide.   Dissolve  a  bit  of 
sodium  iodide  as  large  as  a  grain  of  wheat  in  2  cc.  of  water  in 
a  test  tube,    (h)  What  is  the  color  of  the  solution?   Now  add 
to  the  solution  i  cc.  of  chlorine  water,    (i)  Account  for  the 
change  in  the  color  of  the  solution. 

4.  Action  of  iodine  on  starch.    Place  two  test  tubes  in  your 
test-tube  rack  and  label  them  No.  i  and  No.  2.   Half  fill  each 
with  starch  solution.   To  No.  i  add  a  few  drops  of  the  aqueous 
solution  of  iodine  reserved  in  experiment  2,  above.   To  No.  2 
add  a  bit  of  sodium  iodide  about  as  large  as  a  grain  of  wheat 
and  shake  the  mixture  until  the  iodide  dissolves.    (;')  Does 
free  iodine  change  the  color  of  starch  solution?*  (k)   Does 

[84] 


FIG.  72.  Preparing  iodine  by 

the  action  of  sulf  uric  acid  on 

a  mixture  of  sodium  iodide 

and  sulfuric  acid 


iodine  in  combination  change  the  color  of  starch  solution? 
Now  add  i  or  2  cc.  of  chlorine  water  to  tube  No.  2  and  shake 
the  mixture.  (/)  Note  and  account  for  the  results.  (If  the 
color  is  very  deep,  pour  out  a  portion  of  the  liquid  and  replace 
it  with  water.) 

5.  (m)  Summarize  the  properties  and  conduct  of  iodine. 

EXERCISE  49 

THE  COMPOUNDS  OF  THE  HALOGENS   (CHLORINE, 
BROMINE,  AND  IODINE)  WITH  HYDROGEN 

Apparatus.    3  test  tubes. 

Materials.  A  solution  of  sulfuric  acid  prepared  by  adding  (care) 
10  cc.  of  the  concentrated  acid  to  2  cc.  of  water;  i  g.  each  of  sodium 
chloride,  sodium  bromide,  and  sodium  iodide;  3  strips  of  blue  litmus 
paper. 

1.  Place  three  test  tubes  in  your  rack  and  label  them  No.  i, 
No.  2,  No.  3,  respectively.    In  No.  i  put  about  i  g.  of  sodium 
chloride,  in  No.  2  a  like  weight  of  sodium  bromide,  and  in 
No.  3  the  same  weight  of  sodium  iodide.   Now  add  to  No  i, 
drop  by  drop,   2  cc.  of  the  sulfuric  acid,  shaking  the  tube 
after  each  addition  and  waiting  for  the  reaction  to  subside. 
Note  that  a  colorless  gas  is  evolved  which  escapes  from  the 
tube  and  attracts  moisture  from  the  air,  forming  a  mist  (this 
is  more  noticeable  if  you  blow  gently  across  the  mouth  of  the 
tube}.    Moisten  a  strip  of  blue  litmus  paper  with  water  and 
hold  it  at  the  mouth  of  the  tube,    (a)  What  is  the  gas  evolved 
(Exercise  29)  ?    (b)  Write  the  equation  for  its  formation. 

2.  In  a  similar  way  add  2  cc.  of  the  sulfuric  acid  to  tube 
No.   2,  as  directed  above.    Note  the  colorless   gas   evolved 
(blow  gently  across  the  mouth  of  the  tube),    (c)  Test  it  with 
blue  litmus  paper  as  before,    (d)  What  is  the  gas?    (e)  Write 
the  equation  for  its  formation.    (/)  Is  a  colored  substance  also 
formed  in  the  reaction?    (g)  Recalling  that  hydrogen  bromide 
is  unstable,  account  for  the  formation  of  the  colored  substance. 

[85] 


3.  Now  add  2  cc.  of  the  sulfuric  acid  to  tube  No.  3.   (h)  Is  a 
colorless  gas  formed  as  before  (blow  across  the  mouth  of  the 
tube)  ?    (i)  Test  with  blue  litmus  paper  as  directed  above. 
(j)  What  is  the  colored  product  formed  (hydrogen  iodide  is 
even  less  stable  than  hydrogen  bromide)  ?    (k)  Do  you  recog- 
nize (odor)  any  gas  evolved  in  addition  to  hydrogen  iodide? 
(/)  Account  for  the  formation  of  this  gas,  recalling  that  hydro- 
gen iodide  is  unstable  and  that  nascent  hydrogen  is  a  strong 
reducing  agent. 

4.  (m)  Why  is  the  litmus  paper  used  in  the  above  experi- 
ments moistened  before  using?    (n)   Contrast  in  properties 
hydrogen  chloride,  hydrogen  bromide,  and  hydrogen  iodide. 
(o)  Give  the  names  of  the  aqueous  solution  of  each. 

EXERCISE  50 

THE  SALTS  OF  THE  BINARY  ACIDS  OF  THE  HALOGENS: 
CHLORIDES,  BROMIDES,  AND  IODIDES 

Apparatus.    3  test  tubes ;  mortar  and  pestle. 

Materials,  i  g.  each  of  sodium  chloride,  sodium  bromide,  sodium 
iodide;  i  cc.  silver  nitrate  solution  (R.S.);  nitric  acid;  10  cc.  chlorine 
water  (R.S.);  5  cc.  carbon  tetrachloride;  2g.  manganese  dioxide;  sul- 
furic acid. 

1.  Test  for  chlorides,  bromides,  and  iodides.  Place  three  test 
tubes  in  your  rack  and  label  them  No.  i,  No.  2,  and  No.  3, 
respectively.  Put  into  No.  i  a  bit  of  sodium  chloride  as  large 
as  a  grain  of  wheat,  in  No.  2  put  a  like  amount  of  sodium  bro- 
mide, and  in  No.  3  the  same  amount  of  sodium  iodide.  Now 
add  to  each  tube  about  2  cc.  of  water  and  shake  the  mixtures 
until  the  solids  are  dissolved;  then  add  to  each  tube  2  or  3 
drops  of  silver  nitrate  solution,  (a)  Describe  the  results. 
(b)  Write  the  reaction  in  each  case.  Add  to  each  tube  4  or  5 
drops  of  nitric  acid  and  shake  the  mixture,  (c)  Does  the  pre- 
cipitate dissolve?  The  reaction  with  silver  nitrate  serves  to 
detect  the  presence  of  chlorides,  bromides,  and  iodides. 

[86] 


2.  How  to  distinguish  between  chlorides,  bromides,  and 
iodides.   Arrange  three  test  tubes  as  directed  in  experiment  i, 
above.   Thoroughly  mix  in  a  mortar  a  bit  of  sodium  chloride 
as  large  as  a  pea  with  an  equal  amount  of  manganese  dioxide 
and  transfer  the  mixture  to  tube  No.  i.    Put  in  tube  No.  2  a 
similar  mixture  of  sodium  bromide  and  manganese  dioxide, 
and  in  tube  No.  3  a  like  mixture  of  sodium  iodide  and  man- 
ganese dioxide.    Set  the  tubes  in  the  hood  and  add  to  each 
4  or  5  drops  of  sulfuric  acid.    If  the  reaction  is  not  marked, 
heat  the  tubes  gently,    (d)  How  can  you  distinguish  between 
chlorides,  bromides,  and  iodides? 

Again  arrange  three  clean  test  tubes  in  your  rack.  Prepare 
in  No.  i  a  solution  of  sodium  chloride  as  directed  in  experi- 
ment i,  above;  in  a  similar  way  prepare  in  tubes  No.  2  and 
No.  3  a  solution  of  sodium  bromide  and  of  sodium  iodide  re- 
spectively. Now  add  to  each  solution  about  i  cc.  of  carbon 
tetrachloride  (note  that  the  tetrachloride  does  not  mix  with  the 
water  but  sinks  to  the  bottom  of  the  tube).  Shake  the  mix- 
tures and  then  set  them  in  the  rack  for  a  few  seconds  until 
the  liquids  separate,  (e)  Note  the  color  of  the  carbon  tetra- 
chloride. Now  add  to  each  tube  about  2  cc.  of  chlorine  water 
and  again  shake  the  contents  of  the  tubes  vigorously ;  then  set 
the  tubes  aside  until  the  liquids  again  separate.  (/)  Now  note 
the  color  of  the  carbon  tetrachloride.  (g)  What  is  the  function 
of  the  chlorine  water?  (h)  Describe  in  detail  the  method  for 
distinguishing  between  chlorides,  bromides,  and  iodides  by 
using  chlorine  water  and  carbon  tetrachloride.  Note  that  the 
silver  nitrate  test  will  enable  you  to  tell  whether  or  not  a  given 
compound  belongs  to  the  group  composed  of  chlorides,  bro- 
mides, and  iodides.  Either  the  test  with  manganese  dioxide 
and  sulfuric  acid  or  with  chlorine  water  and  carbon  tetra- 
chloride will  enable  you  to  distinguish  the  three  members  of 
the  group  from  each  other. 

3.  Obtain  from  your  instructor  an  unknown  sample  of  a 
chloride,  bromide,  or  iodide  and  determine  its  identity. 

[87] 


EXERCISE  51 

THE  PROPERTIES  OF  GASOLINE  AND  KEROSENE; 
ACETYLENE 

Apparatus.  2  test  tubes ;  250-00.  wide-mouthed  bottle ;  medicine 
dropper  or  pipette. 

Materials.  Filter  paper ;  i  cc.  cottonseed  oil ;  wooden  splint ;  3  cc. 
gasoline ;  3  cc.  kerosene ;  a  bit  of  calcium  carbide  (CaC2)  the  size  of 
a  bean. 

1.  Properties  of  gasoline  and  kerosene.  (Gasoline  is  easily 
ignited.  The  student  must  keep  it  away  from  all  flames.) 
Draw  two  circles  (as  large  as  a  silver  dollar)  on  a  filter  paper. 
Mark  one  circle  G  and  the  other  K.  Within  the  circle  marked 
G  place  i  drop  of  gasoline  (use  a  medicine  dropper)  and  within 
the  circle  marked  K  place  i  drop  of  kerosene.  Suspend  the 
paper  so  that  both  sides  of  the  paper  are  exposed  to  the  air, 
and  note  from  time  to  time,  (a)  Which  product  is  the  more 
volatile  ? 

Pour  about  2  cc.  of  gasoline  into  a  test  tube  and  add  2  drops 
of  some  fat  or  oil  such  as  cottonseed  oil.  (b)  Does  the  oil  dis- 
solve in  the  gasoline?  (c)  What  use  does  this  suggest  for 
gasoline  ? 

Pour  about  i  cc.  of  gasoline  into  a  test  tube  and  add  a  like 
volume  of  water,  (d)  Do  the  two  liquids  mix?  (e)  Which  is 
the  heavier?  (/)  Repeat,  using  kerosene,  (g)  Can  burning 
gasoline  be  extinguished  by  water  ? 

Gently  warm  a  250-0:.  wide-mouthed  bottle  by  rotating  it 
above  a  flame.  Add  to  the  warm  bottle  2  or  3  drops  of  gasoline, 
place  the  palm  of  your  hand  tightly  over  the  mouth  of  the 
bottle,  and  shake  the  bottle  vigorously.  Remove  your  hand 
and,  standing  at  arm's  length,  bring  a  lighted  splint  to  the 
mouth  of  the  bottle,  (h)  Record  the  results,  (i)  What  ap- 
plication of  gasoline  is  based  on  the  property  noted  in  this 
experiment  ? 

[88] 


2.  Acetylene.  Nearly  fill  a  test  tube  with  water  and  drop 
into  it  a  small  piece  of  calcium  carbide.  (;)  What  is  the 
gas  evolved?  (k)  Write  the  equation  for  its  preparation. 
(/)  Ascertain  by  holding  a  lighted  splint  at  the  mouth  of  the 
tube  whether  the  gas  is  inflammable. 

EXERCISE  52 

SOME  DERIVATIVES  OF  METHANE :  CHLOROFORM, 
CARBON  TETRACHLORIDE,  IODOFORM 

Apparatus.    Test  tubes;  beaker  (ioo-cc.);  glass  rod. 

Materials.  2  cc.  chloroform  ;  i  cc.  cottonseed  oil ;  3  cc.  carbon  tetra- 
chloride ;  splint ;  2  cc.  alcohol ;  5  cc.  iodine  solution  (R.  S.) ;  sodium 
hydroxide  solution. 

1.  Chloroform,  (a)  Note  its  odor.   Add  about  i  cc.  of  chloro- 
form to  an  equal  volume  of  water,    (b)   Does  it  mix  with 
Water?    (c)  Is  it  heavier  or  lighter  than  water?   Add  2  drops 
of  an  oil  (such  as  cottonseed  oil)  to  2  cc.  of  chloroform,  (d)  Is 
chloroform  a  good  solvent  for  oils?    (e)  Determine  whether 
chloroform  will  burn  by  dipping  the  end  of  a  glass  rod  into  the 
liquid  and  then  bringing  it  in  contact  with  the  edge  of  a  flame 
for  an  instant. 

2.  Carbon  tetrachloride.    (/)    Repeat   all   the  experiments 
listed  under  chloroform,  substituting  carbon  tetrachloride  for 
the  chloroform,    (g)  Pour  about  2  cc.  of  carbon  tetrachloride 
into  a  small  beaker  (100  cc.)  and  heat  it  until  the  vapor  of  the 
liquid  fills  the  beaker ;  then  thrust  a  lighted  splint  into  the 
vapor.    Mention  two  important  commercial  uses  of  carbon 
tetrachloride  illustrated  in  the  above  experiments,    (h)  What 
is  the  great  advantage  of  carbon  tetrachloride  over  gasoline  as 
an  agent  in  cleaning  clothes? 

3.  lodoform.    Pour  about  2  cc.  of  alcohol  into  a  small  beaker 
and  add  5  cc.  of  iodine  solution.    Now  add  to  the  mixture  a 
solution  of  sodium  hydroxide,  drop  by  drop,  stirring  the  liquid 
after  each  addition.    Continue  until  the  color  of  the  iodine  is 

[891 


completely  discharged.  Note  the  formation  of  a  yellowish  solid 
(iodoform).  (i)  Note  its  odor.  (;')  What  is  the  chief  use  of 
the  compound  ? 

EXERCISE  53 
A  STUDY  OF  FLAMES 

Apparatus.    Wire  gauze;  porcelain  dish;  blowpipe. 

Materials.  Charcoal  (size  of  a  bean)  ;  5  cc.  limewater  (R.  S.)  ;  can- 
dle; wooden  splint;  piece  of  charcoal  8x3x2  cm.;  bit  of  lead  as  large 
as  a  pea;  2  g.  lead  oxide. 

1.  Parts  of  a  flame,  (a)  Note  and  account  for  the  difference 
between  the  combustion  of  a  wooden  splint  and  that  of  a  piece 
of  charcoal,  (b) 
What  are  the  con- 
ditions necessary 
for  the  production 
of  a  flame?  (c) 
Light  a  candle  and 
place  it  so  that  the 
flame  is  against  a 
black  background 
and  is  not  dis- 
turbed by  air 
drafts;  then  note 
the  different  cones 
in  the  flame.  Test 
the  relative  tem- 

.     ....       FIG.  73.  Using  a  blowpipe  to  form  reducing  and  cxi- 

peratures    of   dif-  dizing  flames 

ferent  parts  of  the 

flame  by  means  of  narrow  strips  of  splints,  (d)  Draw  a  dia- 
gram in  your  notebook  showing  the  different  parts  of  the  flame 
and  noting  the  color  of  each  of  the  cones.  Extinguish  the  candle 
flame  and  hold  a  lighted  splint  2  or  3  cm.  from  the  wick  in  the 
little  column  of  smoke  rising  from  the  wick,  (e)  Note  and 
account  for  the  result. 

[90] 


2.  Products  of  combustion.    (/)  What  two  elements  consti- 
tute the  main  percentage  of  ordinary  fuels?  (g)  What  products 
form  when  these  elements  burn  in  air  or  oxygen?    (h)  Devise 
and  describe  simple  experiments  to  show  the  presence  of  these 
products  in  the  gases  evolved  by  the  burning  candle. 

3.  Kindling  temperature,   (i)  What  is  meant  by  the  kindling 
temperature  of  gases?   When  a  lamp  is  first  lighted,  a  film  of 
liquid  often  spreads  over  the  chimney  for  an  instant,    (j)  Ex- 
plain.   Press  a  piece  of  wire  gauze  halfway  down  on  a  Bunsen 
flame.    Notice   that   the 

flame  does  not  extend 
above  the  gauze,  (k)  Is 
this  due  to  the  absence 
there  of  gases  that  are 
combustible  (test  for  their 
presence  by  means  of  a 
lighted  splint)  ? 

Turn  off  the  gas,  then 
turn  it  on  and  ignite  it 
over  a  piece  of  wire  gauze 
held  horizontally  4  or  5  cm.  above  the  top  of  the  burner. 
(/)  Note  the  results  and  explain,  (m)  How  does  the  miner's 
safety  lamp  prevent  explosions? 

4.  Luminous  flames.    Hold  a  porcelain  dish  in  a  small  lumi- 
nous Bunsen  flame,    (n)  Account  for  the  deposition  of  carbon. 
(0)  Does  the  nonluminous  flame  deposit  carbon?    (p)  To  what 
is  the  luminosity  of  the  flame  due?    Recall  the  experiment  on 
the  Bunsen  flame  in  Exercise  2. 

5.  Use  of  the  blowpipe.    Close  the  openings  on  the  Bunsen 
burner  so  as  to  secure  a  luminous  flame ;  then  lower  the  flame 
until  about  3  cm.  high.    Now  place  the  end  of  the  blowpipe 
just  inside  the  lower  part  of  the  flame  and  blow  gently  through 
the  blowpipe  (Figs.  73,  74).   The  flame  will  be  bent  over,  as 
shown  in  the  figures.    Practice  this  until  you  can  produce  a 
continuous  and  steady  flame  by  forcing  a  portion  of  the  exhaled 

[911 


Reducing  flame 

Oxidizing  flame 


FIG.  74.  Showing  the  location  of  the  reducing 
and  the  oxidizing  flames 


air  through  the  blowpipe  by  constant  tension  of  the  muscles  of 
the  cheek.  It  will  be  noted  that  the  flame  so  produced  consists 
of  two  cones,  one  inside  the  other.  The  inner  cone  being  rich  in 
carbon  monoxide  is  a  reducing  flame ;  the  outer  cone  being  very 
hot  and  in  contact  with  oxygen  of  the  air  is  an  oxidizing  flame. 

Make  a  shallow  cavity  near  the  end  of  a  piece  of  charcoal 
and  place  in  this  a  bit  of  lead.  Now  heat  this  in  the  oxidizing 
flame  as  shown  in  Fig.  73.  The  lead  is  changed  to  lead  oxide 
(PbO),  a  yellow  solid  which  collects  about  the  cavity  on  the 
surface  of  the  charcoal. 

Clean  the  charcoal  and  place  in  the  cavity  a  paste  of  lead 
oxide  (PbO)  made  by  adding  a  little  water  to  2  g.  of  the  oxide 
and  mixing.  Now  heat  this  in  the  reducing  flame  (hold  the 
charcoal  so  that  the  oxide  is  in  the  tip  of  the  inner  cone).  The 
oxide  will  be  gradually  reduced  to  a  bead  of  metallic  lead. 

NOTE.  Save  the  charcoal  for  future  experiments,  but  before  return- 
ing it  to  your  desk,  pour  some  water  into  the  cavity  to  extinguish  any 
fire  which  may  be  smoldering  in  the  interior  of  the  charcoal;  otherwise 
it  is  a  source  of  danger  from  fire.  Indeed,  to  avoid  all  danger  from 
fire,  it  would  be  better  if  the  charcoal  used  by  the  students  were  kept 
in  some  noncombustible  vessel  such  as  a  crock  or  pneumatic  trough. 

EXERCISE  54 
SOME  COAL-TAR  COMPOUNDS  (OPTIONAL) 

Apparatus.  Test  tubes;  25o-cc.  flask,  fitted  with  stopper  and  glass 
tube  as  shown  in  Fig.  75;  glass  rod;  25o-cc.  beaker. 

Materials.  10  cc.  benzene;  2  drops  of  cottonseed  oil;  nitric  acid; 
sulfuric  acid ;  i  or  2  cc.  aniline ;  5  g.  bleaching  powder. 

NOTE.  The  student  should  first  start  experiment  2;  while  the  liquids 
are  being  heated,  experiment  i  can  be  performed. 

Caution.  Benzene  is  very  inflammable  and  must  be  kept 
away  from  all  flames. 

1.  Benzene  (C6H6).  (a)  What  is  the  distinction  in  com- 
position between  benzene  and  benzine?  (b)  Note  the  odor 

[92] 


of  benzene,  (c)  Add  about  i  cc.  of  benzene  to  an  equal  vol- 
ume of  water,  (d)  Does  it  mix  with  water?  (e)  Is  it  heavier 
or  lighter  than  water?  Add  2  drops  of  cottonseed  oil  to  2  cc. 
of  benzene.  (/)  Is  it  a  good  solvent  for  oils?  (g)  Dip  the 
end  of  a  glass  rod  into  the  liquid  and  touch  it  to  the  edge  of 
a  flame  for  an  instant,  (h)  Is  benzene 
inflammable  ? 

2.  Nitrobenzene  (C6H5NO2).  Pour  8cc. 
of  nitric  acid  into  a  250-0:.  flask  and  add 
(care),  drop  by  drop,  with  constant  motion 
of  the  flask  so  as  to  mix  the  liquids,  10  cc. 
of  sulfuric  acid,  cooling  the  mixture  all  the 
time  by  allowing  cold  water  to  run  on  the 
outside  of  the  flask.  When  the  mixture  is 
cold  add,  drop  by  drop,  6  cc.  of  benzene, 
mixing  the  liquids  thoroughly  after  each 
addition  of  benzene.  Now  connect  the  flask 
with  a  cork  and  glass  tube  (Fig.  75)  and 
heat  the  mixture  very  gently  (not  to  boiling 
point).  The  tube  acts  as  an  air  condenser  to 
prevent  the  escape  of  the  benzene.  During 
the  heating  mix  the  liquids  in  the  flask 
every  few  minutes  by  gently  rotating  the 
flask.  Continue  the  heating  for  about  half 
an  hour ;  then  disconnect  the  flask,  cool  it, 
and  pour  the  contents  into  a  250-0:.  beaker  filled  with  cold 
water.  Stir  the  mixture  with  a  glass  rod,  and  after  the  oil 
(nitrobenzene)  settles  to  the  bottom,  decant  the  clear  liquid, 
leaving  the  oil  in  the  beaker.  Fill  the  beaker  again  with  water, 
stir,  and  decant  as  before.  In  this  way  the  acids  are  washed 
away,  leaving  the  nitrobenzene  pure. 

The  nitric  acid  reacts  with  the  benzene  to  form  nitrobenzene 
and  water  as  shown  in  the  following  equation: 


FIG.  75.  Preparing  ni- 
trobenzene by  heating 
benzene  and  a  mixture 
of  nitric  and  sulfuric 
acids 


C«EL  +  HNO, 


C6H5N02 


[93] 


The  sulfuric  acid  assists  by  absorbing  the  water  formed  in 
the  reaction,  (i)  Record  the  color  of  the  nitrobenzene;  also 
its  odor  (poisonous  if  too  much  is  inhaled)  and  density  as  com- 
pared with  water. 

The  above  is  the  first  step  in  the  commercial  production  of 
aniline,  which  is  the  source  material  for  many  dyes.  The  nitro- 
benzene is  converted  into  aniline  by  reducing  it : 

C6H5NO2  +  3H2 >•  C6H5NH2  (aniline)  +  2H2O 

3.  Aniline  (C6H5NH2).  (j)  Note  the  properties  (color,  odor, 
solubility  in  water)  of  aniline.  Add  one  drop  of  aniline  to 
100  cc.  of  water  in  a  flask  and  shake  the  mixture.  The  aniline 
soon  dissolves  in  the  water.  Place  about  i  g.  of  bleaching  pow- 
der in  a  beaker,  add  2  5  cc.  of  water,  and  stir  the  mixture  thor- 
oughly with  a  glass  rod ;  then  filter  and  add  i  or  2  cc.  of  the 
filtrate  to  the  aniline  solution  and  mix  thoroughly.  If  no  change 
in  color  takes  place,  add  more  of  the  filtrate,  2  or  3  cc.  at  a  time. 
(k)  Record  your  results.  This  reaction  serves  as  a  delicate 
color  test  for  aniline. 

EXERCISE  55 
THE  SUGARS 

Apparatus.    3  test  tubes;  2  small  beakers;  stirring-rod. 

Materials.  10  cc.  each  of  Fehling's  solution  A  and  B  (R.  S.)  ;  5  cc. 
commercial  glucose  or  corn  sirup;  2  g.  sucrose;  i  g.  each  of  sweets 
such  as  candy,  honey,  molasses ;  2  or  3  g.  sodium  carbonate  dissolved 
in  as  little  water  as  possible;  red  litmus  paper;  hydrochloric  acid. 

1.  The  action  of  Fehling's  solution  on  dextrose  (grape  sugar). 
The  most  common  test  for  dextrose  is  the  reaction  with 
Fehling's  solution.  The  latter  is  prepared  by  mixing  equal 
volumes  of  solutions  A  and  B,  prepared  as  directed  in  the 
Appendix. 

Pour  into  a  test  tube  about  i  cc.  each  of  solutions  A  and  B. 
When  thoroughly  mixed,  the  resulting  solution  should  be  deep 
blue,  but  perfectly  clear.  Heat  the  blue  solution  nearly  to 

[94] 


boiling,  add  i  or  2  drops  of  commercial  glucose  (corn  sirup, 
procured  at  any  grocery,  will  do  as  well),  and  continue  the 
heating  for  a  few  moments.  The  copper  sulfate  in  the  solu- 
tion is  reduced  to  cuprous  oxide  by  the  dextrose,  and  this 
separates  in  the  form  of  a  red  or  yellow  solid.  Levulose  will 
act  in  the  same  way.  Dissolve  samples  of  candy,  honey,  and 
molasses  in  a  little  water  and  test  for  the  presence  of  dextrose 
and  levulose  in  these  sweets,  (a)  Note  the  results  obtained. 

2.  The  action  of  Fehling's  solution  on  sucrose.  Try  the 
action  of  Fehling's  solution  on  sucrose  as  in  i.  (b)  Note  the 
results  obtained.  Now  dissolve  about  0.5  g.  of  the  sugar  in  5  cc. 
of  water  in  a  test  tube  and  add  3  or  4  drops  of  concentrated  hy- 
drochloric acid.  Place  the  tube  in  a  beaker  of  boiling  water  and 
leave  it  there  for  about  five  minutes;  then  pour  the  solution 
into  a  small  beaker,  cool  the  solution,  and  neutralize  the  acid 
present  by  adding  a  concentrated  solution  of  sodium  carbonate 
until  the  resulting  mixture  is  just  alkaline  (test  with  red  litmus 
paper).  Now  test  this  with  Fehling's  solution  as  in  i.  (c)  Ac- 
count for  the  result. 

EXERCISE  56 

THE  COMPOSITION  OF  MILK 

Apparatus.    Evaporating-dish ;  stirring-rods. 

Materials.  150-00.  milk;  acetic  acid  (R. S.);  Fehling's  solution 
(R.S.) ;  one  half  junket  tablet  (these  tablets  may  ordinarily  be  obtained 
from  any  druggist;  they  may  always  be  had  at  very  little  cost  by  ad- 
dressing Chr.  Hansen's  Laboratory,  Little  Falls,  New  York). 

1.  Determination  of  the  solids  and  the  water  in  milk.  Place 
a  short  stirring-rod  in  a  small  evaporating-dish  and  accurately 
weigh  the  two  together,  recording  the  weights  in  the  table  below. 
Introduce  about  20  cc.  of  milk  and  again  weigh.  Now  evapo- 
rate the  milk  to  dryness  (Fig.  58),  occasionally  stirring  it  with 
the  rod  to  break  the  scum  (protein  matter),  which  would  other- 
wise retard  evaporation.  While  the  evaporation  is  taking  place 
proceed  with  2. 

[95] 


When  the  residue  is  perfectly  dry,  carefully  wipe  the  outside 
of  the  dish,  cool,  reweigh,  and  record  the  weight.  From  your 
results  calculate  the  percentages  of  solids  and  water  in  the 
milk.  Pure  milk  contains  not  less  than  12  per  cent  of  solid 
matter.  Consult  your  state  laws  as  to  the  minimum  amount 
of  solid  matter  allowed  in  milk  sold  in  the  state. 

Weight  of  evaporating-dish  and  stirring-rod    ...  g. 

Weight  of  evaporating-dish,  stirring-rod,  and  milk  .  g. 

Weight  of  milk  taken  (calculate) g. 

Weight  of  dish,  rod,  and  residue  left  after  evaporation  g. 

Weight  of  solids  in  milk  taken  (calculate)  ....  g. 

Percentage  of  solids  in  milk  taken  (calculate)  ...  % 
Percentage  of  water  in  milk  taken  (ioo%—      %  of 

solids) v  •    ?, % 

2.  The  separation  of  the  protein  (casein)  in  milk  by  rennin. 
Dissolve  about  one  half  of  a  so-called  "junket"  tablet  in  a  little 
cold  water  and  add  the  solution  to  about  100  cc.  of  sweet  milk 
heated  until  just  lukewarm.    Stir  the  solution,  then  cool  it  and 
set  it  aside  until  the  rennin  contained  in  the  tablet  causes  the 
separation  of  the  curd  (casein).    Break  up  the  curd  by  stirring, 
and  filter  it.   Save  the  filtrate  for  further  tests,    (a)  For  what 
is  the  curd  used? 

3.  The  separation  of  the  protein  in  milk  by  acid.   Add  i  or  2 
drops  of  acetic  acid  (vinegar  will  serve)  to  5  cc.  of  milk,  mix 
thoroughly,  and  set  aside.    The  presence  of  an  acid  causes  the 
casein  to  separate  (hence  the  appearance  of  sour  milk). 

4.  The  separation  of  the  lactose  in  milk.    Evaporate  the  fil- 
trate reserved  in  2,  refiltering  it  if  more  solid  matter  separates. 
(b)  Taste  the  residue,    (c)   Test  it  with  Fehling's  solution 
(Exercise  55). 

5.  (d)  Enumerate  the  different  constituents  you  have  found 
in  the  milk. 


96] 


EXERCISE  57 

THE  DETERMINATION  OF  THE  FAT  PRESENT  IN  MILK 

(OPTIONAL) 

Apparatus.  A  Babcock  milk  tester  (Fig.  76)  consisting  of  a  centrifu- 
gal machine  (^4),  2  test  bottles  (B),  pipette  made  to  deliver  17.6  cc.  (C), 
and  a  small  cylinder  (D)  graduated  to  hold  17.5  cc. ;  beaker  (5oo-cc.). 

Materials.  Pint  bottle  of  the  milk  to  be  tested ;  sulfuric  acid  (con- 
centrated, density  1.82  to  1.83;  the  ordinary  commercial  sulfuric  acid 
serves  well) ;  100  cc.  boiling  water. 

NOTE.  The  most  highly  prized  content  of  milk  is  the  fat.  This  is 
sometimes  in  part  removed  by  unscrupulous  dealers  and  sold  for  cream 
or  used  in  making  butter;  or  water  is  added  to  the  milk,  which  reduces 


-q 

( 

—  IWfcri 

10 

- 

0 

- 

# 

jal 

B 

)(- 

B 

} 

E 


D 


FIG.  76.  The  apparatus  used  in  determining  the  percentage  of  fat  in  milk  by 
the  Babcock  method 


the  percentage  of  fat.  To  protect  the  public,  laws  have  been  passed  by^ 
the  different  states  as  well  as  by  the  Federal  government  fixing  the  mini- 
mum percentage  of  fat  in  milk  sold  on  the  market.  Normal  milk  con- 
tains from  3  per  cent  to  5  per  cent  of  fat,  averaging  about  4  per  cent. 
The  laws  of  most  of  the  states  require  not  less  than  3  per  cent.  The 
method  ordinarily  used  for  determining  the  fat  in  milk  was  devised  by 
Professor  Babcock  and  is  known  as  the  Babcock  test. 

[97] 


In  this  test  a  definite  volume  of  milk  is  mixed  with  sulfuric  acid  in 
a  so-called  test  bottle,  the  neck  of  which  is  graduated  so  as  to  read  per- 
centages of  fat.  The  heat  evolved  in  the  mixing  of  the  two  is  sufficient 
to  destroy  the  casein  of  the  milk  but  leaves  the  fat.  The  bottle  is  then 
whirled  in  a  centrifugal  machine;  in  this  way  the  fat,  being  light,  is 
finally  drawn  into  the  neck  of  the 
bottle,  and  the  amount  present  can 
be  read  off. 

In  order  to  balance  the  centrif- 
ugal machine  properly  it  is  neces- 
sary to  have  two  test  bottles,  equally 
filled,  opposite  each  other.  Two 
different  samples  can  therefore  be 
tested  at  the  same  time,  or  in  case 
only  one  sample  is  to  be  tested, 
both  bottles  can  be  filled  with 
samples  of  the  same  milk,  and  the 
results  compared. 

1.  Introducing  the  milk  and 
acid  into  the  test  bottles.   The 
milk  to  be  tested  is  first  mixed 
thoroughly  by  pouring  it  from 
the  bottle  into  a  beaker,  and 
vice  versa,  three  or  four  times. 
By  means  of  the  pipette,  exactly 

i7.6cc.  of  the  milk  is  then  transferred  to  each  of  the  test 
bottles,  as  shown  in  Fig.  77.  Next  fill  the  17.5  cc.  cylinder  with 
the  sulfuric  acid  to  the  graduation  mark ;  then,  holding  each  of 
the  test  bottles  in  turn  at  an  angle,  slowly  pour  the  17.50:. 
of  acid  down  the  wall  of  the  bottle,  turning  the  bottle  at  the 
same  time  so  as  to  wash  down  the  neck  of  the  bottle  any  milk 
.adhering  to  the  wall.  The  acid,  being  heavier  than  the  milk, 
sinks  to  the  bottom  of  the  bottle. 

2.  Mixing  the  acid  and  milk.    The  milk  and  acid  in  each 
bottle  is  now  carefully  mixed  by  gently  rotating  the  bottle. 
Considerable  heat  is  developed  in  the  process.    Care  must  be 
taken  not  to  throw  the  mixture  into  the  neck  of  the  bottle. 

[98] 


FIG.  77.  Showing  the  proper  way 

of  adding  milk  to  the  test  bottle, 

determining  the  fat  in  milk 


Continue  until  the  acid  and  milk  are  thoroughly  mixed  (the 
layer  of  acid  at  the  bottom  of  the  bottle  must  entirely  dis- 
appear with  the  formation  of  a  dark  homogeneous  liquid). 

3.  Separating  the  fat.   The  two  bottles  are  at  once  placed 
opposite  each  other  in  the  containers  (E,  E,  Fig.  76)  of  the 
centrifugal  machine,  and  the  bottles  whirled  for  four  or  five 
minutes.    The  rate  of  whirling  varies  with  the  machine  (di- 
rections are  given  with  each  machine).   The  machine  is  then 
stopped,  and  hot  (nearly  boiling)  water  is  added  to  each  bottle 
until  the  liquid  reaches  nearly  to  the  bottom  of  the  narrow  neck. 
The  bottles  are  again  whirled  for  one  minute.   The  machine  is 
again  stopped,  and  hot  water  is  added  until  the  lower  part  of 
the  fat  column  reaches  about  to  the  i  or  2  per  cent  mark  on  the 
neck.    The  bottles  are  again  whirled  one  or  two  minutes.   The 
fat  should  now  be  clear  and  should  come  within  the  percentage 
marks  on  the  bottle.   By  taking  the  lower  and  upper  readings 
of  the  fat  column  and  subtracting,  the  percentage  of  the  fat  in 
the  milk  is  obtained. 

4.  (a)  Give  the  results  of  your  tests. 


EXERCISE  58 

A  STUDY  OF  STARCH 

Apparatus.    Microscope;  2oo-cc.  beaker;  stirring-rod;  3  test  tubes. 

Materials,  i  cc.  of  iodine  solution  (R.  S.) ;  o.i  g.  flour  (largely  wheat 
starch) ;  10  g.  corn  starch ;  piece  of  bread ;  hydrochloric  acid ;  3  g. 
sodium  carbonate  dissolved  in  a  little  water ;  red  litmus  paper ;  starch 
solution  (R.S.);  Fehling's  solution  (R.S.) 

1.  Microscopic  appearance.    Examine  under  the  microscope 
the  appearance  of  starch  from  different  sources  (corn,  wheat) 
when  magnified,    (a)  Draw  diagrams  of  the  starch  granules. 

2.  Actions  of  acids  on  starch.    Try  the  action  of  starch  solu- 
tion on  Fehling's  solution  (as  in  Exercise  55).    (b)  Note  your 
results. 

[99] 


Add  2  cc.  of  hydrochloric  acid  to  50  cc.  of  starch  solution  in  a 
beaker  and  boil  the  contents  gently  for  thirty  minutes,  allowing 
the  solution  to  concentrate  to  about  2  5  cc.  Cool  the  liquid, 
neutralize  with  sodium  carbonate,  and  again  test  the  solution 
with  Fehling's  solution,  (c)  Again  note  the  results  and  explain. 

3.  Test  for  starch.    Recall  the  action  of  iodine  on  starch 
(Exercise  48).   This  constitutes  a  good  test  for  starch. 

Test  different  foods  (such  as  bread,  potatoes,  and  corn  meal) 
for  starch.  To  do  this,  boil  about  5  g.  of  the  food  with  100  cc. 
of  water,  stirring  the  mass  thoroughly  so  as  to  break  it  into 
small  pieces ;  then  filter  it  and  cool  the  filtrate.  Now  stir  the 
filtrate  with  a  glass  rod,  the  end  of  which  is  first  dipped  into  a 
solution  of  iodine. 

4.  The  action  of  heat  on  starch.    Place  3  or  4  g.  of  starch  in 
a  test  tube  and  heat  slightly  for  ten  or  fifteen  minutes,  stirring 
it  often  with  a  glass  rod  and  regulating  the  heat  so  as  not  to 
burn  the  starch  (the  same  results  may  be  obtained  by  heating 
a  piece  of  bread  in  an  oven  until  it  is  dry  and  crisp).   How 
does  the  product  differ  in  taste  from  the  original  starch  ?    The 
heat  changes  a  part  of  the  starch  into  an  isomeric  compound 
known  as  dextrin,   (d)  For  what  is  dextrin  used? 

EXERCISE  59 
TEXTILE  FIBERS;  PAPER 

Apparatus.  4  small  beakers  or  test  tubes ;  stirring-rod ;  (evaporating- 
dish;  2  large  beakers;  microscope;  2.5o-cc.  flask  fitted  with  stopper  and 
tube  as  in  Fig.  75). 

Materials.  3  strips  each  of  uncolored  cotton,  pure  wool  (the  kind 
known  as  nun's  veiling  is  apt  to  be  pure),  silk,  and  linen  cloth  (3  cm. 
xiscm.);  50  cc.  sodium  hydroxide  solution;  strip  of  filter  paper; 
hydrochloric  acid  ;  sulfuric  acid ;  (3  strips  of  filter  paper  2  cm.  x  10  cm. ; 
ammonium  hydroxide ;  sodium  hydroxide,  desk  solution ;  piece  of  old 
white  cotton  cloth). 

1.  Effect  of  heat  upon  textile  fibers.  Ignite  the  end  of  a  strip 
of  cotton  cloth  in  a  Bunsen  flame ;  then  withdraw  it  from  the 

[100] 


flame,    (a)  Note  the  bdor;o?  tlie 3&ftapg ;# dtii;   (6)  Does  the 
cloth  when  ignited  continue  to  burn? 

Repeat,  using  strips  of  wool,  silk,  and  linen,  (c)  Can  you 
distinguish  in  this  way  between  vegetable  fibers  (cotton,  linen) 
and  animal  fibers  (wool,  silk)  ? 

2.  How  to  distinguish  between  vegetable  fibers   (cotton, 
linen)  and  animal  fibers  (wool,  silk).    Place  a  strip  of  each 
kind  of  cloth  in  small  beakers,  cover  the  cloth  with  sodium 
hydroxide  solution,  and  boil  the  liquid  for  ten  minutes,  replac- 
ing the  water  as  it  evaporates  (the  cloth  must  always  be  com- 
pletely covered  with  the  liquid) ;  then  set  the  beakers  aside 
until  cool,    (d)  Record  the  results.    Cloth  sold  as  woolen  often 
contains  considerable  cotton,    (e)  How  could  you  detect  this 
adulteration?    It  will  be  of  interest  to  test  samples  of  woolen 
cloth  to  see  if  they  are  pure. 

3.  How  to  distinguish  between  the  animal  fibers.    Immerse 
strips  of  silk  and  wool  in  concentrated  hydrochloric  acid  and 
(/)  note  the  change  after  they  have  stood  a  few  minutes. 

4.  How  to  distinguish  between  the  vegetable  fibers.  Immerse 
strips  of  cotton  and  linen  in  concentrated  sulfuric  acid  for  two 
minutes,    (g)  Record  the  results. 

5.  Miscroscopic  appearance  of  textile  fibers.    Examine  the 
appearance  of  each  kind  of  fiber  when  magnified,    (h)  Com- 
pare with  the  diagrams  given  in  your  text. 

6.  Parchment  paper.    Pour  20  cc.  of  sulfuric  acid  (care}  slowly  and 
with  constant  stirring  into  a  beaker  containing  locc.  of  water.    Pour 
the  solution  into  an  evaporating-dish  and  allow  to  cool.    Draw  strips  of 
filter  paper  slowly  through  the  acid  and  then  immerse  them  in  a  large 
beaker  of  water.    Finally,  wash  the  strips  in  a  large  beaker  of  water 
containing  5  or  10  drops  of  ammonium  hydroxide. 

(i)  When  the  strips  are  dry  compare  their  properties  with  those  of 
the  untreated  paper. 

7.  Paper  pulp.    Cut  into  short,  thin  strips  a  piece  of  old  white  cot- 
ton cloth.    Introduce  about  2  g.  of  these  strips  into  a  25o-cc.  flask  and 
add  about  20  cc.  of  a  solution  of  the  sodium  hydroxide  on  your  desk. 
Insert  in  the  mouth  of  the  flask  a  stopper  through  which  passes  a  glass 

noii 


tube  (Fig.  75) ;  £h%i?  JiQ^tVtae  flr.s-k  until  t'he  liquid  just  boils,  for  two 
or  three  hours.  When  cool  decant  the  liquid  and  wash  the  residue 
several  times  with  water  until  it  is  free  from  sodium  hydroxide.  This 
residue  is  a  form  of  pulp  used  in  making  fine  writing  paper.  By  beat- 
ing the  pulp  until  it  is  very  fine,  straining  it  through  a  fine-wire  sieve, 
and  then  pressing  it  with  a  hot  iron,  it  is  possible  to  obtain  a  coarse 
kind  of  handmade  paper. 


EXERCISE  60 

THE  PREPARATION  AND  PROPERTIES  OF  COMMMON 
ALCOHOL 

Apparatus.  One  2ooo-cc.  flask  (or  bottle)  connected  with  tube  and 
bottle,  as  shown  in  Fig.  78 ;  test  tube ;  stirring-rod ;  apparatus  shown 
in  Fig.  57;  evaporating-dish ;  distilling  apparatus  (Fig.  41). 

Materials.  200  g.  molasses  or  corn  sirup  ;  cake  of  yeast ;  50  cc.  lime- 
water  (R.  S.)  ;  25  cc.  alcohol ;  small  amounts  (size  of  a  pea)  of  sugar, 
starch,  salt;  i  cc.  cottonseed  oil;  5  cc.  iodine  solution  (R.  S.)  ;  sodium 
hydroxide  solution;  soda  lime  (a  mixture  of  sodium  hydroxide  and  cal- 
cium hydroxide)  sufficient  to  fill  tube  C  (Fig.  78). 

1.  Preparation  of  alcohol.  (It  is  suggested  that  this  experi- 
ment be  performed  by  the  instructor  or  by  students  selected 
by  the  instructor ;  after 
the  alcohol  is  generated, 
the  liquid  may  be  di- 
vided among  the  differ- 
ent members  of  the  class 
who  will  then  test  for 
the  alcohol  as  directed 
in  4.) 

Dissolve  about  200  g. 
of  ordinary  molasses  (or 
sirup)  in  2000  cc.  of 
water  in  the  flask  A 
(Fig.  78).  Grind  a  cake  of  yeast  with  a  little  water  and  add  it 
to  the  solution  in  A.  Connect  the  flask  as  shown  in  the  figure 
(the  bottle  B  contains  limewater  and  the  tube  C  contains  soda 

[102] 


FIG.  78.  Preparation  of  alcohol  by  the  fermen- 
tation of  molasses 


lime).  Set  the  apparatus  aside  in  a  warm  place  (30°  is  best) 
for  two  or  three  days.  Note  that  a  gas  is  evolved  in  A  and 
bubbles  through  the  limewater  in  B.  (a)  Write  the  equation 
for  the  reaction  taking  place  in  A  and  B. 

2.  Properties  of  alcohol,  (b)  Determine  the  boiling  point  of 
alcohol,  using  the  apparatus  shown  in  Fig.  57.   (c)  Pour  a  few 
drops  of  alcohol  into  an  evaporating-dish,  ignite,  and  note  the 
characteristics  of  the  flame. 

(d)  Determine  whether  alcohol  is  a  good  solvent  for  sugar, 
salt,  starch,  and  oils,  such  as  cottonseed  oil.  (e)  Does  alcohol 
mix  with  water.  (/)  Contrast  the  properties  of  alcohol  with 
those  of  gasoline. 

3.  Test  for  alcohol.  Pour  2  cc.  of  alcohol  into  a  test  tube  and 
add  to  this  5  cc.  of  iodine  solution.   Now  add  a  solution  of 
sodium  hydroxide,  one  drop  at  a  time  (mix  after  the  addition 
of  each  drop),  until  the  iodine  color  vanishes;   then  warm 
gently  and  set  aside  for  a  few  minutes.   A  yellow  precipitate 
of  iodoform  (Exercise  52),  of  characteristic  odor,  forms.    (If 
the  amount  of  alcohol  present  is  small,  the  iodoform  may  not 
separate,  but  its  presence  will  be  revealed  by  its  odor.) 

4.  When  the  fermentation  of  the  sugar  in  A   (Fig.  78)  is 
complete,  divide  the  liquid  so  that  each  student  or  group  of 
students  will  have  from  150  to  200  cc.    Pour  the  liquid  into 
a  flask  and  distill  (Fig.  41)  3  or  4  cc.    (g)  Dip  the  end  of  a 
glass  rod  in  the  distillate  and  touch  it  to  the  edge  of  a  flame. 
Test  the  remainder  of  the  filtrate  for  alcohol  as  in  3  above. 

EXERCISE  61 
THE  COMPOSITION  OF  FLOUR 

Apparatus.    Beaker  (500-0:.);  porcelain  crucible;  pipestem  triangle. 
Materials.    50 g.  flour;  iodine  solution  (R.S.)  for  testing  for  starch; 
cheesecloth  (from  12  to  15  cm.  square). 

1.  Weigh  out  about  25  g.  of  flour  and  mix  with  just  enough 
water  to  form  a  stiff  dough,  working  it  in  the  hands  until  it 

[103] 


becomes  smooth  and  elastic.  Place  the  dough  on  a  piece  of 
cheesecloth,  then  fold  the  cloth  about  the  dough  and  tie  it  with 
a  string  so  as  to  form  a  little  bag.  Nearly  fill  your  largest 
beaker  with  water,  immerse  the  bag  and  contents  in  the  water, 
and  repeatedly  squeeze  the  bag  between  the  fingers.  Note  that 
the  water  becomes  cloudy.  Retain  a  portion  of  this  cloudy 
liquid  for  tests  in  2. 

Continue  the  washing  in  fresh  portions  of  water  until  the 
resulting  wash  water  remains  clear.  The  action  may  be  has- 
tened by  working  the  dough  in  a  small  stream  of  running  water. 
The  residue  is  the  nitrogenous  constituent  of  flour  and  is  known 
as  gluten.  Burn  a  portion  of  the  gluten  and  (a)  note  the  odor. 

2.  Test  separate  portions  of  the  cloudy  liquid  reserved  in 
i   for  starch     (Exercise   58)    and   for  sugar    (Exercise   55). 
(b)  Record  your  results. 

3.  Place  about  o.i  g.  of  flour  in  a  porcelain  crucible,  heat  it 
with  the  Bunsen  flame  (Fig.  32),  and  gradually  increase  the 
heat  until  only  a  white  residue  remains.  The  residue  is  mineral 
matter. 

4.  (c)  Enumerate  the  different  constituents  you  have  found 
in  the  flour. 

EXERCISE  62 
THE  ACTION  OF  PRESERVATIVES  (OPTIONAL) 

Apparatus.  Two  250-00.  bottles  or  beakers;  soo-cc.  beaker;  2  test 
tubes. 

Materials.  300  cc.  sweet  milk  ;  drop  of  formalin ;  10  cc.  hydrochloric 
acid  to  which  is  added  one  drop  of  a  solution  of  ferric  chloride  (R.  S.)  ; 
i  g.  sodium  benzoate ;  loog.  tomato  catchup. 

1.  The  action  of  formaldehyde  on  milk.  Thoroughly  clean 
two  small  bottles  with  hot  water  and  half  fill  each  with  sweet 
milk.  Add  to  the  milk  in  one  of  the  bottles  one  drop  of  formalin 
and  mix  thoroughly.  Now  pour  about  3  cc.  of  milk  from  each 
of  the  two  bottles  into  separate  test  tubes ;  add  to  each  an  equal 
volume  of  the  hydrochloric  acid  to  which  has  been  added  one 

[104] 


drop  of  ferric  chloride  solution  (the  ordinary  commercial 
hydrochloric  acid  serves  the  purpose  well,  since  it  usually  con- 
tains a  trace  of  ferric  chloride  as  an  impurity). 

Mix  the  contents  of  each  of  the  tubes  thoroughly  and  set 
them  in  a  beaker  of  boiling  water  (Fig.  79).  Note  any  change 
in  color,  (a)  How  can  you  detect  the  pres- 
ence of  formaldehyde  in  milk?  (b)  What 
is  the  distinction  between  formalin  and 
formaldehyde  ? 

Set  the  two  bottles  containing  the  re- 
mainder of  the  milk  aside  and  examine 
from  day  to  day,  noting  when  the  milk  in 
each  becomes  sour. 

2.  The  action  of  sodium  benzoate  on 
tomato  catchup.  The  student  may  like- 
wise study  the  action  of  sodium  benzoate 
in  preventing  the  fermentation  of  tomato 
catchup.  Pour  the  catchup  in  small  beak- 
ers, or  bottles,  adding  the  benzoate  to 
one  of  the  samples,  mixing  it  thoroughly. 
Allow  the  samples  to  stand  exposed  to  the  air  and  (c)  note 
which  is  the  first  to  sour  (odor  and  taste).  In  such  cases  the 
weight  of  the  preservative  added  to  the  catchup  should  be 
from  o.i  to  0.2  per  cent  of  the  weight  of  the  catchup. 


FIG.  79.  Heating  liquids 

in  test  tubes,  by  placing 

them  in  boiling  water 


105 


EXERCISE  63 

ACETIC  ACID  :  A  STUDY  OF  VINEGAR 

Apparatus.  Evaporating-dish  ;  beaker  (ioo-cc.);  graduated  test  tube; 
(burette;  graduated  pipette  ;  250-00.  flask). 

Materials.  5  cc.  acetic  acid  (R.  S.)  ;  30  cc.  different  kinds  of  vinegar, 
such  as  cider  and  distilled  or  white  ;  i  cc.  phenolphthalein  solution 
(R.  S.);  ordinary  desk  sodium  hydroxide  solution;  (sodium  hydroxide 
solution  of  known  strength.  A  so-called  normal  solution  serves  the  pur- 
pose well.  This  contains  40  g.  of  the  hydroxide  in  1  1.  of  solution  and 
may  be  purchased  at  any  supply  house). 

1.  The  properties  of  acetic  acid.  Pour  about  5  cc.  of  acetic 
acid  into  a  small  beaker,  (a)  Note  its  odor,  (b)  Test  it  with 
blue  litmus  paper,  (c)  Add  to  the  acid  about  50  cc.  of  water, 
mix  thoroughly,  and  taste  one  drop.  Now  add  to  the  dilute 
acid  2  or  3  drops  of  phenolphthalein  solution  and  then  add, 
drop  by  drop,  the  sodium  hydroxide  solution  on  your  desk, 
stirring  the  liquid  thoroughly  after  each  addition.  Continue 
until  the  solution  is  neutral  (see  Exercise  31).  The  reaction  is 
expressed  by  the  following  equation  : 


NaOH  +  H  •  C2H3O2  —  ^NaC2H3O2  +  H2O 

(acetic  acid)  (sodium  acetate) 

The  sodium  acetate  formed  is  a  white  solid  which  may  be  re- 
covered by  evaporating  the  water,  (d)  To  what  class  of  com- 
pounds does  sodium  acetate  belong  ? 

2.  Determination  of  the  solids  in  vinegar  (quantitative). 
In  this  experiment  different  students  should  use  different  kinds 
of  vinegar  and  compare  results.  Weigh  a  small  evaporating^ 
dish  and  record  the  weight  in  Table  I.  Pour  into  the  dish  a 
definite  volume  (say  25  cc.)  of  vinegar.  Evaporate  the  vinegar 
to  complete  dry  ness  (Fig.  58)  ;  then  carefully  wipe  the  bottom 
of  the  dish  with  a  dry  towel  and  again  weigh,  recording  the 
weight  in  the  table. 

[106] 


From  your  results  calculate  the  amount  of  solid  matter  in 
TOO  cc.  of  the  vinegar.  Pure  cider  vinegar  should  contain  not 
less  than  1.6  g.  of  solids  in  100  cc.  of  vinegar  (this  is  the  limit 
fixed  by  the  Federal  government  as  well  as  by  the  statutes  of 
many  of  the  states).  Distilled  vinegar  or  white  vinegar,  on 
the  other  hand,  contains  only  a  trace  of  solid  matter,  (e)  Note 
the  odor  and  taste  of  the  solid  matter  obtained  from  the  vine- 
gar. The  solids  from  pure  cider  vinegar  should  have  an  odor 
and  taste  suggestive  of  baked  apples.  The  character  of  the 
solids  varies  according  to  the  source  of  the  vinegar. 

TABLE  I 

Weight  of  evaporating-dish g. 

Weight  of  evaporating-dish  and  residue    .     .     ,     .    « .  g. 
Weight  of  solids  (residue)  in  volume  of  vinegar  taken 

(calculate)       g. 

Weight  of  solids  in  100  cc.  of  vinegar  (calculate)  .  g. 

3.  Determination  of  the  acidity  of  vinegar.  The  Federal  statutes 
require  that  all  vinegar  sold  shall  contain  not  less  than  4  g.  of  acetic 
acid  in  100  cc.  of  vinegar.  The  amount  of  acid  present  may  be  deter- 
mined in  the  following  way  : 

Prepare  or  purchase  a  dilute  solution  of  sodium  hydroxide  of  known 
strength.  If  a  normal  solution  of  sodium  hydroxide  is  used,  dilute  ex- 
actly 10  cc.  of  it  to  100  cc.  with  water  and  mix  thoroughly,  i  cc.  of  this 
diluted  solution  then  contains  exactly  0.004  g.  of  sodium  hydroxide, 
which  is  a  convenient  strength  to  use.  By  means  of  a  graduated  pipette 
(or  burette)  run  into  a  25o-cc.  flask  exactly  5  cc.  of  the  vinegar  to  be 
tested  and  dilute  with  about  50  cc.  of  water.  Add  to  this  2  drops  of 
phenolphthalein  solution.  Next  fill  a  burette  with  the  sodium  hydroxide 
solution  and  run  this,  drop  by  drop,  into  the  flask  until  the  solution  is 
neutral,  following  the  directions  given  in  Exercise  32.  The  liquids  must 
be  mixed  after  each  addition  of  the  hydroxide  by  gently  rotating  the 
flask.  From  the  volume  of  sodium  hydroxide  solution  used,  calculate 
the  weight  of  the  hydroxide  required  to  neutralize  the  acetic  acid  in 
the  vinegar  (see  equation  in  i,  above).  From  this  result  calculate  the 
weight  of  acetic  acid  present  in  the  5  cc.  of  vinegar  taken;  then  from 
this,  the  weight  of  acetic  acid  in  100  cc.  of  vinegar.  (/)  Record  the 
results  obtained  in  Table  II. 

[107] 


TABLE  II 

Weight  of  sodium  hydroxide  in  i  cc.  of  the  solution  used  g. 

No.  of  cc.  of  hydroxide  solution  used  to  neutralize  the 

acid  in  5  cc.  of  vinegar 

Weight  of  sodium  hydroxide  required  to  neutralize  the 

acid  in  5  cc.  of  vinegar  (calculate)     ......  g. 

Weight  of  acetic  acid  in  5  cc.  of  vinegar  (calculate)  .     .  g. 

Weight  of  acetic  acid  in  100  cc.  of  vinegar  (calculate)  g. 

EXERCISE  64 
ESTERS :  FATS  AND  OILS 

Apparatus.    25o-cc.  flask ;  test  tubes  ;  stirring-rod. 

Materials.  10  cc.  glacial  acetic  acid  (R.S.);  sulfuric  acid;  10  cc. 
alcohol ;  2  or  3  cc.  fats  and  oils,  such  as  cottonseed  oil,  olive  oil,  lard, 
butter,  and  cheese;  5  cc.  carbon  tetrachloride  (R.S.)  ;  piece  of  white 
writing  paper. 

1.  Preparation  of  a  simple  ester  (ethyl  acetate),  (a)  What  is 
an  ester  ?  Ethyl  acetate  is  an  example  of  a  simple  ester.  It  is 
a  colorless  liquid  boiling  at  78°  and  has  a  rather  pleasant  odor. 
It  is  derived  from  acetic  acid  (H  •  C2H3O2)  by  replacing  one 
atom  of  hydrogen  by  the  univalent  radical  known  as  ethyl 
(C2H5).  This  change  is  easily  brought  about  by  heating  acetic 
acid  with  alcohol,  thus : 

H  •  C2H302  +  C2H5OH — >-  C2H5  •  C2H3O2  +  H2O 

(acetic  acid)  (alcohol)  (ethyl  acetate) 

Prepare  ethyl  acetate  in  the  following  way:  Into  a  250-0:. 
flask  pour  locc.  of  glacial  acetic  acid  and  add  to  it  (care), 
drop  by  drop  and  with  constant  mixing,  5  cc.  of  sulfuric  acid 
and  then  3  cc.  of  alcohol.  Apply  a  gentle  heat  to  the  mixture. 
Ethyl  acetate  is  formed  and  may  be  recognized  by  its  charac- 
teristic fragrant  odor.  Do  not  mistake  the  odor  of  alcohol  for 
that  of  ethyl  acetate  (warm  a  few  drops  of  alcohol  in  a  test 
tube  and  note  the  difference  between  its  odor  and  that  of  the 
ethyl  acetate).  The  sulfuric  acid  assists  in  the  reaction  be- 
cause of  its  affinity  for  water. 

[108] 


2.  Composition  and  properties  of  fats  and  oils,  (b)  To  what 
class  of  compounds  do  the  fats  belong?  (c)  Give  the  name  and 
valence  of  the  hydrocarbon  radical  present  in  fats,  (d)  Give 
the  formulas  and  names  of  the  three  chief  constituents  of  fats 
and  oils,  (e)  What  is  the  main  difference  in  composition  be- 
tween solid  fats  and  oils?  (/)  Test  the  solubility  of  typical 
fats  and  oils  in  water,  alcohol,  and  carbon  tetrachloride  as  in 
Exercise  52.  There  is  no  simple  accurate  test  for  oils  and  fats, 
but  the  following,  based  on  the  fact  that  fats  are  nonvolatile, 
will  serve  fairly  well:  In  different  places  on  a  piece  of  white 
writing  paper  rub  one  drop  of  various  fats,  such  as  cottonseed 
oil,  cream,  melted  lard,  and  butter,  (g)  Note  the  appearance 
of  the  resulting  spots  when  held  in  front  of  a  light,  (h)  Heat 
the  paper  slightly  to  see  if  the  spots  will  disappear,  (i)  Test 
different  samples  of  cheese  (warm  a  small  piece  and  rub  it  on 
paper)  and  milk  for  fats  by  this  method. 

EXERCISE  65 

METHODS  FOR  DISTINGUISHING  BETWEEN  BUTTER 
AND  OLEOMARGARINE  (OPTIONAL) 

Apparatus.  Iron  spoon  ;  wooden  splint  (match  stick  will  do) ;  loo-cc. 
beaker. 

Materials.  10  g.  each  of  fresh  butter  and  oleomargarine  (also  process 
butter,  if  available) ;  50  cc.  sweet  milk. 

1.  Foam  test.    Over  a  small  flame  heat  gently  in  a  spoon  2  or 
3  g.  of  the  sample  (butter  or  oleomargarine).    Under  these  con- 
ditions butter  will  melt  without  sputtering  and  with  the  forma- 
tion of  much  foam  on  the  surface,  while  oleomargarine  will 
sputter  and  give  but  little  foam.    Process  butter,  or  renovated 
butter  (rancid  butter  which  has  been  purified  by  melting  the 
fat,  skimming  it  off  the  surface,  and  churning  it  with  milk 
under  certain  conditions),  acts  like  oleomargarine,    (a)  Record 
the  results  of  your  tests. 

2.  Sweet-milk  test.    Pour  50  cc.  of  sweet  milk  into  a  small 
beaker  and  heat  nearly  to  boiling.   To  the  hot  milk  add  4 

[109] 


or  5  g.  of  the  sample  and  stir  it  with  a  wooden  splint  until  the 
fat  is  melted,  then  place  the  beaker  in  ice  water  and  continue 
the  stirring  until  the  fat  solidifies.  Under  these  conditions 
butter  will  solidify  in  the  form  of  granules  which  mix  with  the 
milk.  Oleomargarine,  on  the  other  hand,  will  collect  in  a 
single  mass  so  that  it  can  be  removed  from  the  milk  in  one 
lump  with  the  wooden  stirrer.  (b)  Record  your  results. 

EXERCISE  66 
PROTEINS 

Apparatus.  Evaporating-dish ;  test  tubes ;  glass  rod ;  loo-cc.  beaker. 
Materials.   Small  portion  of  white  of  an  egg ;  25  cc.  milk ;  nitric  acid ; 
ammonium  hydroxide  ;  i  g.  flour ;  small  strip  of  woolen  cloth. 

1.  Proteins.  Place  a  small  portion  of  the  white  of  an  egg 
(protein)  in  a  test  tube  and  heat  it  by  dipping  the  tube  into 
boiling  water,  (a)  Record  the  results. 

Pour  20  cc.  of  milk  into  an  evaporating-dish  and  heat  gently. 
Note  the  scum  that  collects  on  the  surface.  Remove  the  scum 
with  a  glass  rod,  collect  other  portions  of  the  scum  in  the  same 
way,  and  save  for  further  tests. 

Transfer  a  portion  of  the  coagulated  white  of  egg  to  a  small 
beaker  and  moisten  it  with  2  or  3  drops  of  nitric  acid,  (b)  Note 
any  change,  (c)  Wash  the  egg  free  from  acid  with  repeated 
portions  of  water,  then  moisten  it  with  ammonium  hydroxide 
and  note  change  in  color.  These  color  changes  serve  as  a  test 
for  protein  matter. 

In  the  same  way  test  for  the  presence  of  protein  in  the  scum 
which  separated  when  the  milk  was  heated,  (d)  Test  other 
substances  for  protein,  such  as  flour,  woolen  cloth,  a  clipping 
of  a  finger  nail,  recording  the  results.  Nitric  acid  stains  the 
skin  yellow ;  (e)  suggest  an  explanation  for  the  change  in  color. 

Burn  a  small  bit  of  different  kinds  of  protein,  such  as  egg, 
hair,  the  scum  of  milk  (note  the  odor).  (/)  Record  your  results. 

[110] 


EXERCISE  67 

PHOSPHORUS  AND  ITS  COMPOUNDS 

Apparatus.  250-00.  wide-mouthed  bottle;  deflagrating-spoon ;  glass 
plate  ;  small  beaker ;  forceps. 

Materials.  Phosphorus  half  the  size  of  a  pea  (to  be  secured  from  your 
instructor  when  needed) ;  litmus  paper  (red  and  blue) ;  10  cc.  ammonium 
molybdate  solution  (R.  S.) ;  ammonium  hydroxide ;  nitric  acid. 

1.  (Read  carefully  the  precautions  given  for  handling  phos- 
phorus in  paragraph  3,  Exercise  8.)  Cover  the  bottom  of  a  wide  • 
mouthed  bottle   (250-00.)    with  water  to  a  depth  of  about 
i  cm.    By  means  of  your  forceps  place  the  piece  of  phosphorus 
on  a  deflagrating-spoon  and  ignite  it  by  touching  it  with  a  hot 
wire.    Quickly  lower  the  phosphorus  into  the  bottle  and  cover 
the  mouth  of  the  bottle  with  a  glass  plate.   When  the  phos- 
phorus ceases  to  burn,  withdraw  the  spoon  and  hold  it  in  the 
Bunsen  flame  for  a  few  seconds  (to  insure  that  all  the  phos- 
phorus is  burned  off  the  spoon).    Cover  the  bottle  with  the 
palm  of  your  hand  and  shake  the  water  so  that  the  fumes  in 
the  bottle  are  dissolved,    (a)   Test  the  solution  with  litmus 
paper,    (b)  What  is  present  in  the  water? 

2.  Pour  the  solution  in  the  bottle  into  a  small  beaker,  add  2 
or  3  cc.  of  nitric  acid,  and  boil  the  solution  until  about  half  of 
it  evaporates.    The  nitric  acid  oxidizes  to  phosphoric  acid  all 
the  phosphorus  compounds  present.   Add  a  few  drops  of  the 
solution  to  10  cc.  of  a  solution  of  ammonium  molybdate  and 
warm  gently,    (c)  Note  the  result  (the  compound  formed  has 
a  very  complex  composition).   Add  ammonium  hydroxide  to 
the  mixture  until  the  liquid  is  alkaline,    (d)  Note  the  result. 
Again  acidify  the  liquid  with  nitric  acid,    (e)  Note  the  result. 
The  formation  of  a  yellow  precipitate  with  ammonium  molyb- 
date, which  precipitate  is  insoluble  in  nitric  acid  and  soluble 
in  ammonium  hydroxide,  serves  as  a  good  test  for  phosphoric 
acid  and  its  salts. 

[mi 


EXERCISE  68 


SOME  COMPOUNDS  OF  ARSENIC 

Apparatus.  Hard-glass  tube  about  12  cm.  long  and  6mm.  internal 
diameter,  to  be  made  by  the  student  by  sealing  the  end  of  a  piece  of 
hard-glass  tubing  (A,  Fig.  80)  ;  magnifying-glass  ;  glass  rod ;  iron  spoon. 

Materials.  0.2  g.  arsenious  oxide  ;  2  or  3  g.  charcoal  in  pieces  about 
the  size  of  a  grain  of  wheat. 

1.  Reduction  of  arsenious  oxide.  Into  the  hard-glass  tube 
introduce  an  amount  of  arsenious  oxide  equal  in  bulk  to 
about  2  grains  of  wheat.  Cover 
this  to  a  depth  of  2  or  3  cm.  with 
the  pieces  of  charcoal  (Fig.  80), 
which  have  been  previously  heated 
to  a  high  temperature  in  an  iron 
spoon  to  expel  any  volatile  mat- 
ter. See  that  the  inner  surface  of 
the  tube  above  the  charcoal  is  per- 
fectly clean. 

Incline  the  tube,  and  heat  (use 
low  flame)  the  upper  portion  of 
the  charcoal  to  a  high  temperature 
(Fig.  80) ;  then,  while  maintaining 
the  charcoal  at  this  temperature, 
gradually  bring  the  lower  part  of 
the  tube  also  into  the  flame.  The 
upper  part  of  the  tube  must  not 

be  heated.   The  arsenious  oxide  is 

FIG.  80.  Reducing  arsenious  ox- 
Changed  into  a  vapor,  which  passes  ide  by  hot  charcoal 

over  the  hot  charcoal  and  is  re- 
duced.   The  resulting  vapor  of  arsenic  condenses  on  the  colder 
portion  of  the  tube  just  above  the  charcoal,    (a)  Note  the 
appearance  of  the  arsenic. 

[112] 


2.  Oxidation  of  arsenic.  Cut  the  tube  as  near  the  bottom  as 
possible  and  remove  the  charcoal,  pushing  it  out  at  the  bottom 
of  the  tube  (if  necessary)  by  a  glass  rod;  then,  inclining  the 
tube  at  an  angle  of  about  45°,  apply  a  very  gentle  heat  to  that 
portion  of  it  which  contains  the  deposit  of  arsenic.  The  air 
passing  through  the  tube  oxidizes  the  arsenic,  forming  small 
white  crystals  of  As2O3,  which  are  slowly  deposited  in  the 
colder  portions  of  the  upper  part  of  the  tube,  (b)  Examine 
them  with  a  magnify  ing-glass,  (c)  Note  their  form. 

EXERCISE  69 
A  STUDY  OF  ANTIMONY 

Apparatus.  Blowpipe;  test  tube;  beaker  (200-00.);  stirring-rod; 
hydrogen  sulfide  generator  (Fig.  68). 

Materials.  Piece  of  charcoal  (8  cm.  x  3  cm.  x  2  cm.)  J  2  pieces  of 
antimony  (size  of  a  grain  of  wheat) ;  hydrochloric  acid ;  nitric  acid ; 
strip  of  zinc  half  as  long  as  a  test  tube  and  narrow  enough  to  go  in  a 
test  tube ;  5  g.  ferrous  sulfide. 

1.  Heat  a  bit  of  antimony  on  charcoal  in  the  oxidizing  flame 
(Fig.  73).  The  product  is  Sb2O3. 

2.  Into  a  test  tube  introduce  a  bit  of  antimony  no  larger 
than  a  grain  of  wheat  and  add  about  3  cc.  of  hydrochloric  acid 
and  then  2  or  3  drops  of  nitric  acid.   After  the  metal  has  dis- 
solved, pour  the  solution  into  100  cc.  of  water  in  a  beaker.    If  a 
precipitate  forms,  add  hydrochloric  acid,  drop  by  drop,  with 
constant  stirring,  until  the  precipitate  dissolves.    Half  fill  a 
test  tube  with  the  resulting  solution  and  insert  a  strip  of  zinc, 
as  in  Exercise  35.    (a)  Note  the  results.  Through  the  remain- 
der of  the  solution  pass  a  few  bubbles  of  hydrogen  sulfide 
(Exercise  42).   The  precipitate  is  Sb2S3.    (b)  Note  its  proper- 
ties.  This  is  the  compound  often  used  in  making  red  rubber. 


[113] 


EXERCISE  70 

A  STUDY  OF  BISMUTH 

Apparatus.  Blowpipe;  test  tube;  beaker  (200-00.);  stirring-rod; 
hydrogen  sulfide  generator  (Fig.  68). 

Materials.  2  pieces  of  bismuth  (size  of  a  grain  of  wheat) ;  piece  of 
charcoal  (8  cm.  x  3  cm.  x  2  cm.)  ;  nitric  acid ;  strip  of  zinc  as  in  Exer- 
cise 69  ;  10  g.  ferrous  sulfide  ;  hydrochloric  acid. 

1.  Heat  a  piece  of  bismuth  on  charcoal  in  the  oxidizing  flame 
(Fig.  73).   Bi2O3  is  formed  and  is  deposited  on  the  charcoal. 
(a)  What  is  the  color  of  the  oxide?    (b)  Write  the  equation 
for  its  formation. 

2.  Into  a  test  tube  introduce  a  bit  of  bismuth  and  add  4  or 
5  drops  of  nitric  acid.    If  the  acid  does  not  entirely  dissolve 
the  bismuth,  add  a  few  drops  more.   When  the  bismuth  is 
dissolved,  pour  the  solution  into  100  cc.  of  water  in  a  beaker. 
If  a  precipitate  forms,  add  nitric  acid,  drop  by  drop,  with 
constant  stirring,  until  the  precipitate  dissolves.    Half  fill  a 
test  tube  with  the  solution  and  insert  in  it  a  strip  of  zinc,  as 
in  Exercise  69.    (c)  Note  the  results.  Through  the  remainder 
of  the  solution  pass  a  few  bubbles  of  hydrogen  sulfide.   The 
precipitate  is  Bi2S3.    (d)  Note  its  properties. 


114] 


EXERCISE  71 

COMPOUNDS  OF  SILICON 

Apparatus.  Evaporating-dish ;  stirring-rod. 

Materials.  2  cc.  water  glass  (solution  of  Na2Si03) ;  hydrochloric  acid. 

1.  (a)  Recall  the  formulas  and  names  of  the  important  acids 
of  silicon. 

2.  Place  2  cc.  of  water  glass  (Na2SiO3)  in  an  evaporating- 
dish,  dilute  with  10  cc.  of  water,  and  add  2  or  3  cc.  of  hy- 
drochloric acid.    Note  the  gelatinous  precipitate  of  H2SiO3. 
(b)  Write  the  equation  for  its  formation.    Evaporate  the  mix- 
ture to  dry  ness  (Fig.  24)  and  heat  the  dish  gently  with  the 
bare  flame.  The  silicic  acid  is  decomposed  into  water  and  sili- 
con dioxide,    (c)  Write  the  equation  for  the  reaction.   When 
cool,  add  water,  filter,  and  examine  the  residue,  (d)  What  is  it? 

3.  (e)  Recall  the  action  of  hydrofluoric  acid  on  silica,  writ- 
ing the  equation  for  the  reaction. 

EXERCISE  72 
COMPOUNDS  OF  BORON 

Apparatus.  Fine  platinum  wire  (piece  5  cm.  long  fused  in  glass  tube 
for  handle,  as  shown  in  Fig.  81);  beaker  (ioo-cc.);  stirring-rod;  funnel; 
medicine  dropper. 

Materials.  8  g.  borax ;  sulfuric  acid ;  i  or  2  drops  of  cobalt  nitrate 
solution  (R.S.);  filter  paper. 

1.  (a)  Write  the  names  and  formulas  for  three  important 
compounds  of  boron. 

2.  Make  a  little  loop  on  the  end  of  a  platinum  wire  and  heat 
it  to  redness  in  a  Bunsen  flame;  then  quickly  bring  the  loop 
in  contact  with  some  borax  and  reheat  in  the  tip  of  the  flame 
(Fig.  81).  The  borax  adhering  to  the  loop  will  swell  up  (owing 
to  the  expulsion  of  the  water  of  hydration)  and  finally  form  a 
clear,  glassy  bead,    (b)  Note  the  color  imparted  to  the  flame. 

[115] 


Moisten  the  bead  with  a  drop  of  sulfuric  acid  and  again 
touch  it  to  the  edge  of  the  flame,    (c)  Note  the  result.    This 
serves  as  a  simple  test  for 
borax. 

Moisten  the  bead  with 
a  drop  of  a  solution  of 
cobalt  compound  and  re- 
heat until  the  bead  is 
transparent  when  cold. 
(d)  Note  the  color  of  the 
bead  now  (the  depth  of 
color  depends  upon  the 
amount  of  cobalt  pres- 
ent). This  property  serves 
as  a  simple  test  for  co- 
balt. Some  of  the  other 

metals  likewise  impart  characteristic  colors  to  the  bead,  as  will 
be  explained  later.  This  method  for  detecting  metals  is  known 
as  the  borax-bead  test. 

3.  Dissolve  5  g.  of  borax  in  150:.  of  boiling  water.  Care- 
fully add  to  the  hot  solution  2  or  3  cc.  of  sulfuric  acid  and  stir. 
The  sulfuric  acid  acts  upon  the  borax  as  follows : 

Na2B4O7  +  H2S04  +  5H2O  -r-+  Na2SO4  +  4H3BO3 

The  boric  acid,  H3BO3,  is  insoluble  and  separates  as  a  pre- 
cipitate. Filter  off  the  precipitate ;  then  place  a  bit  of  it  on  the 
loop  of  the  platinum  wire  used  in  2,  above,  and  note  the  color 
it  imparts  to  the  flame,  (e)  Account  for  the  test  for  borax 
given  in  2. 


FIG.  81.  Forming  a  borax  bead 


116 


EXERCISE  73 

COLLOIDS;  EMULSIONS 

Apparatus.   5  test  tubes;  filter;  250-00.  flask;  funnel. 

Materials.  Crystal  of  sodium  thiosulfate;  hydrochloric  acid;  filter 
paper;  about  i5g.  clay  or  garden  soil;  dilute  sulfuric  acid  made  by 
adding  5  drops  of  sulfuric  acid  to  5  cc.  of  water  (care) ;  solution  of 
aluminium  sulfate  (R.  S.);  2  g.  gelatin;  white  of  egg;  5  cc.  kerosene; 
sufficient  thin  shavings  of  white  soap  cut  from  a  bar  of  soap  to  fill  the 
bottom  of  a  test  tube  to  a  depth  of  about  2  cm. 

1.  Dissolve  a  crystal  of  sodium  thiosulfate  about  the  size  of 
a  pea  in  10  cc.  of  water  and  add  i  cc.  of  dilute  hydrochloric 
acid.   Note  the  separation  of  finely  divided  sulfur,  which  is 
formed  in  accordance  with  the  following  equation : 

Na2S203  +  2HC1  — >•  2NaCl  +  H2O  +  SO2  +  S 

{a)  Is  sulfur  soluble  in  water?  Filter  the  mixture,  (b)  Does 
the  sulfur  collect  on  the  filter  paper  ?  Allow  the  liquid  to  stand 
until  the  end  of  the  hour,  (c)  Does  the  sulfur  settle?  (d)  In 
what  state  is  this  sulfur? 

2.  Put  a  lump  of  moist  clay  or  garden  soil  about  the  size  of 
a  small  marble  in  a  250-0:.  flask  and  add  about  100  cc.  of  dis- 
tilled water.  Stopper  the  flask  and  shake  it  vigorously.  Allow 
the  coarse  material  to  settle  for  a  few  minutes.  Pour  off  equal 
portions  of  the  muddy  water  into  three  tubes  marked  respec- 
tively A,  B,  and  C.    Place  tube  A  in  the  test-tube  rack  for 
comparison.    To  tube  B  add  i  or  2  cc.  of  dilute  sulfuric  acid. 
Shake  it  vigorously  and  set  it  beside  tube  A.   To  tube  C  add 
i  or  2  cc.  of  a  solution  of  aluminium  sulfate  and  shake  the  tube 
vigorously ;  then  set  the  tube  in  the  rack  with  the  other  two. 
(e)  Note  the  rate  at  which  suspended  material  in  the  three  tubes 
settles.  Any  electrolyte  will  serve  in  place  of  the  two  reagents 
you  have  used.    (/)  Why  does  a  river  emptying  into  the  ocean 
fill  up  at  its  mouth  ? 

[117] 


3.  Place  in  a  test  tube  a  sufficient  amount  of  small  pieces  of 
dry  gelatin  to  make  a  layer  i  or  2  centimeters  in  depth.   Add 
about  10  cc.  of  water  and  heat  the  water  until  a  clear  liquid  is 
obtained.    Now  immerse  the  tube  in  ice  water  and  note  the 
formation  of  a  gel.  Again  heat  the  substance  until  a  clear  liquid 
is  obtained,  and  then  repeat  the  cooling  process,    (g)  Does  the 
gel  "set"  again? 

4.  Add  a  small  portion  of  the  white  of  an  egg  to  10  cc.  of  cold 
water  and  shake  the  mixture  thoroughly.    The  albumen  of  the 
egg  is  present  in  the  water  in  the  form  of  a  colloid.    Divide  the 
mixture  into  two  equal  portions.    Heat  the  one  portion  to  boil- 
ing,   (h)  Is  the  colloid  precipitated?    To  the  other  portion  add 
2  or  3  drops  of  nitric  acid,    (i)  Note  the  results.    Make  suitable 
test  to  determine  whether  the  coagulation  of  the  albumen  of  the 
egg  is  a  reversible  process. 

5.  Select  two  test  tubes  of  the  same  size.    Place  in  the  bot- 
tom of  one  to  a  depth  of  about  2  cm.  some  thin  pieces  of  white 
soap  cut  from  a  bar  of  soap ;  then  fill  the  tube  about  two  thirds 
full  of  distilled  water.   Warm  the  mixture  and  shake  it  gently 
until  the  soap  dissolves.   Now  fill  the  other  tube  about  two 
thirds  full  of  water.   Next  pour  into  each  tube  about  2  cc.  of 
kerosene.   Note  that  the  oil  floats  on  the  top  of  the  water  in 
each  tube. 

Now  take  a  tube  in  each  hand,  close  the  mouth  of  the  tube 
firmly  with  your  thumb,  and  shake  the  tube  and  contents  vigor- 
ously for  one  minute.  (;)  Note  the  appearance  of  the  liquids 
in  each  tube,  (k)  What  is  such  a  mixture  called?  Set  the 
tubes  aside  five  or  ten  minutes  (or  until  the  next  laboratory 
period)  and  again  examine  the  liquids  in  the  tubes.  (/)  Have 
the  oil  and  water  separated  in  both  tubes?  (m)  What  term  is 
applied  to  a  substance,  such  as  soap,  that  prevents  or  at  least 
delays  the  separation  of  the  liquids  in  an  emulsion  ? 


[118] 


EXERCISE  74 

GENERAL  METHODS  FOR  THE  PREPARATION  OF  THE 
COMPOUNDS  OF  THE  METALS 

Apparatus.    6  test  tubes. 

Materials,  o.ig.  of  each  of  the  following  salts,  dissolved  in  5  cc. 
water:  (i)  calcium  chloride,  (2)  lead  nitrate,  (3)  barium  chloride, 
(4)  ferric  chloride,  (5)  silver  nitrate,  (6)  potassium  iodide  (solutions 
of  any  of  these  on  the  Reagent  Shelf  may  be  used) ;  sodium  carbonate 
solution  (R.S.);  hydrochloric  acid;  sulfuric  acid;  ammonium  hydroxide. 

1.  By  the  direct  union  of  the  elements.    Recall  the  formation 
of  sulfides  of  copper  and  of  iron  (Exercise  40) ;  of  the  chloride 
of  antimony  (Exercise  28).    (a)  .Write  the  equations  for  the 
reactions  involved. 

2.  By  dissolving  a  metal  or  its  hydroxide  in  appropriate 
acids.    Recall  the  formation  of  zinc  sulfate  (Exercise  9) ;  of 
sodium  chloride  (Exercise  31) ;  of  copper  nitrate  (Exercise  37). 
(b)  Write  the  equations  for  the  reactions,    (c)  When  a  metal 
or  its  hydroxide  is  acted  upon  by  an  acid,  what  becomes  of 
the  metal? 

3.  By  acting  upon  a  salt  of  an  acid  with  an  acid  having  a 
higher  boiling  point.    Recall  the  action  of  sulfuric  acid  upon 
sodium  nitrate  (Exercise  37) ;  of  hydrochloric  acid  upon  iron 
sulfide  (Exercise  41) ;  of  sulfuric  acid  upon  fluorides  (Exer- 
cise 46) ;    of  sulfuric  acid   on   chlorides    (Exercise   29) ;    of 
hydrochloric  acid  on  carbonates  (Exercise  33).    (d)  Write  the 
equations  for  each  reaction  and  show  in  what  respects  they  are 
all  similar. 

4.  By  the  decomposition  of  a  compound.    Recall  the  action 
of  heat  upon  potassium  chlorate  (Exercise  7) ;  upon  copper 
nitrate  and  lead  nitrate  (Exercise  38).    (e)  Write  the  equations 
for  the  reactions  involved. 

5.  The  following  compounds  are  insoluble   (see  Appendix, 
Table  of  Solubilities  of  some  of  the  compounds  of  the  metals) : 

[  119  ] 


calcium  carbonate  (CaCO3),  lead  sulfate  (PbSO4),  barium 
carbonate  (BaCO3),  ferric  hydroxide  ( Fe ( OH )3),  silver  chlo- 
ride (AgCl),  lead  iodide  (PbI2).  Prepare  a  small  amount  of 
each  in  a  test  tube.  (/)  Write  the  equation  for  the  reaction 
involved  in  each  case. 

EXERCISE  75 
THE  COMPOUNDS  OF  SODIUM 

Apparatus.  Evaporating-dish  ;  beaker  ;  watch  glass ;  stirring-rod ; 
hard-glass  test  tube;  platinum  wire;  piece  of  cobalt  glass  10  cm.  square; 
magnifying-glass. 

Materials.  5  g.  sodium  carbonate ;  litmus  papers  (red  and  blue) ; 
hydrochloric  acid  ;  sulfuric  acid ;  2  or  3  clear  crystals  of  Glauber's  salt ; 
3  g.  sodium  bicarbonate  ;  i  cc.  limewater  (R.  S.) ;  2  g.  sodium  chloride. 

1.  Recall  experiments  with  sodium  (Exercise  30). 

2.  Dissolve  5  g.  of  sodium  carbonate  in  20  cc.  of  water.   Test 
the  solution  with  red  and  with  blue  litmus  paper,    (a)  Account 
for  the  results. 

Now  convert  the  so- 
dium carbonate  present 
into  common  salt,  (b) 
Describe  the  method  and 
write  the  equation  for  the 
reaction  involved.  Treat 
some  of  the  salt  so  pre- 
pared with  sulfuric  acid. 
(c)  What  gas  is  evolved? 

3.  (d)  Write  the  com- 
plete formula  for  Glau- 
ber's salt.     Select   some 
clear  crystals  of  the  salt, 

place  them  on  a  watch  glass,  and  expose  them  to  the  air  for  one 
hour  or  more,  (e)  Note  the  change.  The  change  is  due  to  the 
fact  that  some  of  the  water  of  hydration  escapes.  Crystals  that 
behave  in  this  way  are  said  to  be  efflorescent. 

[120] 


FIG.  82.  Method  of  making  a  flame  test 


4.  Fill  a  hard-glass  test  tube  about  one  fourth  full  of  sodium 
bicarbonate  and  heat  gently.    Prove  that  carbon  dioxide  is 
evolved  (Fig.  54).   What  liquid  condenses  in  the  colder  part  of 
the  tube?    (/)  Write  equations  for  the  reactions  by  which  so- 
dium carbonate  is  converted  into  the  bicarbonate  and  vice  versa. 

5.  Bend  the  end  of  a  platinum  wire  into  the  form  of  a  small 
loop  and  hold  it  in  the  lower  part  of  the  outer  film  of  the  Bunsen 
flame  (Fig.  82)  until  it  ceases  to  give  any  color  to  the  flame; 
then  dip  it  into  a  solution  of  a  compound  of  sodium  so  that  a 
drop  of  the  solution  is  suspended  in  the  loop ;  then  hold  it  in 
the  Bunsen  flame  as  before,   (g)  Note  the  color,   (h)  Note  the 
appearance  of  the  sodium  flame  when  viewed  through  a  piece 
of  cobalt  glass. 

6.  (i)  Recall  the  action  of  hydrochloric  acid  on  sodium  thio- 
sulfate  (Exercise  73,  i). 

7.  Dissolve  about  2  g.  of  sodium  chloride  in  a  little  water  in 
a  beaker  and  set  the  uncovered  solution  aside  until  the  next 
laboratory  period;  then  (;)  examine  the  shape  of  the  crystals 
with  a  magnifying-glass.    Save  the  crystals  for  reference  in  a 
future  exercise. 

EXERCISE  76 

THE  DETERMINATION  OF  THE  WEIGHT  OF  COMMON  SALT 
OBTAINED  BY  ADDING  HYDROCHLORIC  ACID  TO  A  DEFI- 
NITE WEIGHT  OF  SODIUM  BICARBONATE  (OPTIONAL) 

Apparatus.  Evaporating-dish  and  watch-glass  cover;  medicine  drop- 
per; beaker;  stirring-rod;  balance. 

Materials,    i  g.  sodium  bicarbonate ;  hydrochloric  acid. 

1.  Carefully  weigh  to  o.oi  g.  the  evaporating-dish  and  watch 
glass,  recording  the  weight  in  the  table  below.  Transfer  to  the 
dish  about  i  g.  of  sodium  bicarbonate  and  reweigh,  recording 
the  weights.  Pour  4  or  5  cc.  of  water  on  the  bicarbonate,  and 
place  the  watch  glass  on  the  dish  so  that  only  the  lip  of  the 
dish  remains  uncovered.  Using  a  medicine  dropper  or  pipette, 
drop  down  the  lip  of  the  dish  2  or  3  drops  of  hydrochloric  acid. 

[1211 


Wait  until  the  effervescence  caused  by  the  escape  of  the  carbon 
dioxide  ceases,  then  add  a  few  drops  more  of  the  acid.  Repeat 
until  the  addition  of  the  acid  no  longer  causes  any  effervescence. 
Now  hold  the  watch  glass  in  the  hand  just  above  the  dish,  and 
with  a  little  water  carefully  rinse  back  into  the  dish  the  liquid 
which  has  collected  on  the  undersurface  of  it.  Remove  the 
watch  glass  and  slowly  evaporate  the  solution  (Fig.  58). 

When  the  solution  has  evaporated  nearly  to  dryness,  cover 
the  dish  with  the  watch  glass  and  heat  the  dish  with  the  tip 
of  the  flame.  Continue  the  heating  until  there  is  no  more 
liquid  left  in  the  dish  or  clinging  to  the  undersurface  of  the 
glass.  Then  withdraw  the  heat  and,  after  the  dish  is  cool, 
reweigh  the  dish,  watch  glass,  and  residue  (sodium  chloride). 

From  your  results  calculate  the  amount  of  salt  formed  from 
i  g.  of  the  bicarbonate,  recording  the  weights.  Make  the  cal- 
culations called  for  in  the  following  table : 

Weight  of  evaporating-dish  and  watch  glass  ....  g 
Weight  of  evaporating-dish  and  watch  glass  plus  sodium 

bicarbonate g. 

Weight  of  sodium  bicarbonate  (calculate)    ....  g. 

Weight  of  dish,  watch  glass,  and  sodium  chloride     .     .  g. 

Weight  of  sodium  chloride  (calculate) g- 

Theoretical  weight  of  sodium  chloride  that  can  be  ob- 
tained from  the  weight  of  sodium  bicarbonate  taken  g. 


[122] 


EXERCISE  77 
SOME  COMPOUNDS  OF  POTASSIUM 

Apparatus.  Evaporating-dish ;  stirring-rod;  beaker  (250-00.)  ;  cobalt 
glass;  platinum  wire;  (funnel). 

Materials.  locc.  potassium  hydroxide  solution  (R.S.);  hydrochlo- 
ric acid;  0.5  g.  potassium  chloride;  (iyg.  sodium  nitrate;  i5g.  potas- 
sium chloride;  filter  paper). 

1.  Pour  10  cc.  of  a  solution  of  potassium  hydroxide  into  an 
evaporating-dish,  neutralize  with  hydrochloric  acid,  and  evapo- 
rate to  dryness.    (a)  What  is  the  product?    (b)  By  what  other 
method  have  you  prepared  this  same  compound  in  a  former 
exercise  ? 

2.  Repeat  5,  Exercise  75,  using  a  solution  of  potassium 
chloride,    (c)  Note  the  appearance  of  the  flame,    (d)   Also 
note  the  appearance  of  the  flame  through  a  piece  of  cobalt  glass. 
(e)  What  is  the  difference  in  the  appearance  of  the  sodium 
flame  and  of  the  potassium  flame  when  viewed  through  the 
cobalt  glass?    (/)   How  could  you  detect  both  sodium  and 
potassium  if  they  were  present  in  the  same  solution? 

3.  Preparation  of  potassium  nitrate.   Dissolve  17  g.  of  sodium  nitrate 
in  15  cc.  of  boiling  water ;  also  15  g.  of  potassium  chloride  in  30  cc.  of 
boiling  water.    Mix  the  two  solutions  in  a  small  beaker  and  evaporate 
(stirring  the  mixture)  to  about  20  cc. ;  then  quickly  filter  the  hot  solu- 
tion and  set  the  filtrate  (Filtrate  A)  aside  until  cold. 

The  reaction  between  potassium  chloride  and  sodium  nitrate  is  re- 
versible, and  the  number  of  grams  of  each  of  the  four  compounds  in- 
volved which  dissolve  in  100  g.  of  water  at  15  °  and  100°  is  as  follows  : 

15°    100°  15°    100° 

Sodium  nitrate  .      .     84     180        Potassium  nitrate    .     26     246 
Potassium  chloride  .33       57         Sodium   chloride      .     36      40 

From  a  study  of  these  solubilities  (g)  what  compound  should  you 
expect  would  separate  when  the  hot  solutions  of  sodium  nitrate  and 
potassium  chloride  are  mixed  together?  (h)  Examine  the  crystals  on 

[123] 


the  filter  paper  with  a  magnifying-glass  and  compare  with  7,  Exercise 
75,  to  see  if  your  conclusion  is  correct,  (i)  Taste  the  crystals. 

(;)  What  solid  should  you  expect  would  separate  from  Filtrate  A 
when  it  is  cooled  ?  (&)  Should  you  expect  it  to  be  pure  ?  Filter  off  the 
solid  and  examine  it  with  a  magnifying-glass.  (/)  Can  you  detect  crys- 
tals of  sodium  chloride  in  this  solid  ? 

Dissolve  the  solid  in  as  little  hot  water  as  possible,  cool  the  solution, 
and  again  filter  off  the  solid.  Repeat  until  no  crystals  of  sodium  chloride 
can  be  detected.  Prove  the  identity  of  this  compound  (test  for  potas- 
sium by  the  flame  and  for  a  nitrate  by  2,  Exercise  38). 


EXERCISE  78 
THE  PROPERTIES  OF  AMMONIUM  COMPOUNDS 

Apparatus.    Evaporating-dish  and  watch-glass  cover ;  5  test  tubes. 

Materials.  Ammonium  hydroxide;  hydrochloric  acid;  sodium  hy- 
droxide solution;  litmus  papers  (red  and  blue);  (0.2  g.  ferrous  sulfate 
solution  (R.S.);  ammonium  carbonate  (R.  S.);  barium  chloride  (R.  S.); 
calcium  chloride  (R.S.)). 

1.  Pour  10  cc.  of  ammonium  hydroxide  into  an  evaporating- 
dish,  neutralize  with  hydrochloric  acid,  and  evaporate  just  to 
dryness  (Fig.  58).  (a)  What  is  the  residue?  (b)  Note  its  odor. 

Introduce  about  one  half  of  the  residue  into  a  test  tube,  add 
a  few  drops  of  sodium  hydroxide  solution,  and  heat  gently. 
(c)  Note  the  odor  of  the  evolved  gas  and  its  action  on  a  moist 
strip  of  red  litmus  paper.  All  ammonium  compounds  evolve 
ammonia  when  heated  with  sodium  hydroxide.  This  reaction 
serves  as  a  good  test  for  all  ammonium  compounds. 

Cover  the  evaporating-dish  containing  the  remainder  of  the 
residue  with  a  watch  glass  and  heat  gently  with  a  small  flame. 
Note  that  the  solid  sublimes ;  that  is,  passes  directly  from  the 
solid  form  into  a  vapor  which  condenses  (partly)  on  the  cold 
surface  of  the  watch  glass. 

2.  Add  a  few  drops  of  ammonium  carbonate  solution  to  separate  solu- 
tions of  a  compound  of  barium  and  of  calcium,    (d)  Write  the  equation, 
for  the  reaction  in  each  case. 

[124] 


3.  Add  2  or  3  drops  of  ammonium  sulfide  to  5  cc.  of  a  solution  of 
ferrous  sulfate.  (e)  Note  the  result  and  write  the  equation  for  the  re- 
action that  takes  place.  Repeat,  using  hydrogen  sulfide  solution  in  place 
of  ammonium  sulfide.  (/)  Account  for  the  difference  in  the  reaction 
(consult  Table  of  Solubilities  in  Appendix). 

EXERCISE  79 

DETECTION  OF  COMPOUNDS  OF  THE  ALKALI  METALS 

(OPTIONAL) 

1.  Recall  such  reactions  of  sodium,  potassium,  and  ammo- 
nium compounds;  also  of  carbonates,  sulfates,  nitrates,  chlo- 
rides, bromides,  iodides,  and  phosphates  as  will  serve  to  identify 
them,  and  outline  a  method  of  procedure  for  the  identification 
of  them.  Then  ask  the  instructor  for  unknown  compounds 
falling  within  this  list  and  identify  them,  recording  your  results. 

EXERCISE  80 

THE  PREPARATION  AND  PROPERTIES  OF  SOAP 

Apparatus.  Evaporating-dish ;  large  beaker  ;  stirring-rod  ;  funnel ; 
small  beaker;  4  test  tubes. 

Materials.  10  cc.  alcohol ;  5  g.  cottonseed  oil  (or  lard) ;  i  g.  sodium 
hydroxide  dissolved  in  2  cc.  water;  filter  paper;  hydrochloric  acid;  mag- 
nesium sulfate  (R.S.);  calcium  chloride  solution  (R.S.);  i  g.  sodium 
chloride  dissolved  in  5  cc.  water. 

1.  Add   10  cc.   of  alcohol   to   5  g.  of  cottonseed  oil  in  an 
evaporating-dish.    To  the  resulting  mixture  add  i  g.  of  sodium 
hydroxide  dissolved  in  2  cc.  of  water.    Evaporate  carefully  (use 
small  flame  and  do  not  let  the  tip  touch  the  dish),  stirring  the 
mixture  constantly  until  the  odor  of  alcohol  can  no  longer  be 
detected,    (a)  Write  the  equation  for  the  reaction  on  the  sup- 
position that  the  oil  is  composed  of  olein  only,    (b)  What  is 
the  process  called?    (c)  What  remains  in  the  dish? 

2.  Add  50  cc.  of  distilled  water  to  the  residue  in  the  dish,  stir 
well  for  a  few  minutes,  and  filter,  if  not  clear.    Pour  a  few  drops 
of  the  filtrate  on  your  hands  and  rub  them  (add  more  distilled 

(125  \ 


water  if  necessary)  to  see  if  the  soap  lathers  freely.  Pour  5  cc. 
of  the  filtrate  into  each  of  three  test  tubes.  To  the  first  add 
2  or  3  drops  of  hydrochloric  acid ;  to  the  second  add  a  few  drops 
of  a  solution  of  magnesium  sulfate.  In  like  manner  add  a  few 
drops  of  a  solution  of  calcium  chloride  to  the  third,  (d)  Note 
what  takes  place  in  each  test  tube,  writing  the  equations  for 
each  of  the  reactions,  (e)  Why  do  waters  containing  calcium 
and  magnesium  compounds  (hard  waters)  form  a  precipitate 
(curdle)  with  soap? 

.3.  Add  a  few  drops  of  sodium  chloride  solution  to  10  cc.  of 
the  filtrate  obtained  in  2.  The  soap  in  the  filtrate  is  not  in  solu- 
tion but  in  colloidal  suspension.  The  salt  causes  the  precipita- 
tion of  the  colloid  (soap).  (/)  What  advantage  is  taken  of  this 
reaction  in  the  manufacture  of  soap  ?  • 

4.  Recall  the  effect  of  soap  in  the  formation  of  emulsions 
(Exercise  73).  (g)  What  influence  has  this  property  upon  the 
cleansing  action  of  soap  ? 

EXERCISE  81 
A  STUDY  OF  SOME  OF  THE  COMPOUNDS  OF  CALCIUM 

Apparatus.  Blowpipe  and  charcoal ;  2  test  tubes ;  glass  tubing  for 
blowing  air  through  solution ;  evaporating-dish ;  watch  glass ;  copper 
penny;  piece  of  window  glass;  25o-cc.  flask;  platinum  wire  for  flame; 
test;  (spectroscope). 

Materials.  3  g.  marble;  hydrochloric  acid;  ammonium  carbonate 
(R. S.);  drop  of  cottonseed  oil;  log.  plaster  of  Paris;  log.  commer- 
cial cyanamide;  red  litmus  paper;  o.i  g.  of  the  chloride  or  nitrate  of 
each  -of  the  metals  calcium,  strontium,  barium  (o.i  g.  each  of  sodium 
chloride  and  potassium  chloride). 

1.  Place  one  or  two  small  pieces  of  marble  about  as  large  as 
a  bean  on  a  piece  of  charcoal  and  heat  them  for  five  minutes 
in  the  oxidizing  blowpipe  flame  (Fig.  73).  When  cool  drop  the 
resulting  mass  into  10  cc.  of  cold  water  in  a  test  tube,  (a)  What 
evidence  have  you  of  a  chemical  reaction  taking  place  ?  Shake 
the  tube  for  about  one  minute;  then  filter  and  collect  the  fil- 

[126] 


trate.  Now  gently  blow  exhaled  air  through  the  clear  filtrate. 
(b)  Note  all  the  changes  and  write  the  equations  for  the  three 
reactions  that  have  taken  place. 

2.  Dissolve  i  or  2  g.  of  marble  in  hydrochloric  acid,  (c)  Write 
the  equation  for  the  reaction,    (d)  What  does  the  effervescence 
indicate?    Evaporate  the  solution  to  dryness   (use  the  bare 
flame  and  evaporate  to  complete  dryness).      (e)  What  is  the 
composition  of  the  residue?    Place  a  small  piece  of  it  on  a 
watch  glass  and  expose  it  to  the  air  for  an  hour  or  more  (it  may 
be  left  until  the  next  laboratory  period).    (/)   Account  for 
the  results. 

Dissolve  the  remainder  of  the  residue  in  the  evaporating-dish 
in  5  cc.  of  water  (filter  if  not  clear)  and  add  to  it  a  few  drops 
of  ammonium  carbonate,  (g)  Note  the  results  and  write  the 
equation  for  the  reaction  involved. 

3.  Place  on  a  glass  plate  a  penny  that  has  been  rubbed  with 
a  drop  of  oil.    Pour  over  the  coin  a  thick  paste  made  by  adding 
a  little  water  to  plaster  of  Paris.    Set  the  glass  plate  aside  until 
the  paste  hardens;  then  remove  the  coin  and  (h)  note  the 
results. 

4.  Place  about  one  tablespoonful  of  commercial  cyanamide 
in  a  25o-cc.  flask,  add  50  cc.  of  water,  and  heat  to  boiling. 
Suspend  a  strip  of  red  litmus  paper  just  inside  the  flask  and 
note  any  change  in  color,    (i)  What  do  the  results  indicate? 
(;')  Can  you  detect  the  odor  of  ammonia  in  the  escaping  steam? 
Account  for  its  formation,    (k)  What  is  the  original  source  of 
the  nitrogen  present  in  the  ammonia  that  is  evolved  (recall  the 
method  of  making  the  cyanamide)  ? 

5.  Try  the  flame  tests  for  each  of  the  following  metals  as  in 
5,  Exercise  75   (use  the  chloride  or  nitrate  of  each  metal) : 
calcium,  strontium,  barium.    (/)  Record  the  results  below. 

METAL  COLOR  OF  FLAME 

Calcium      .     .     .     .     .     . 

Strontium  .      .      .,,-.'.      .      . 
Barium 

[1271 


6.  The  use  of  the  spectroscope.  The  spectroscope  is  an  instrument 
which  has  been  of  great  service  in  many  chemical  investigations,  es- 
pecially in  the  discovery  of  new  elements  and  in  detecting  the  presence 
of  known  elements  in  complex  mixtures.  The  principle  involved  in  its 
construction  is  as  follows : 

When  a  beam  of  light  passes  through  a  triangular  prism  of  glass,  it 
is  bent  out  of  its  course  and  emerges  at  a  decided  angle  with  its  original 
direction,  as  shown  in  Fig.  83.  Ordinary  light  is  made  up  of  many 
different  wave  lengths,  and  each  one  is  deflected,  or  refracted,  to  a  dif- 
ferent degree,  so  that  the  various  colors  of  which  the  light  is  composed 
are  spread  out  in  a  series,  the  red 
being  the  least  refracted,  the  violet 
the  most  so.  A  beam  of  white  light 
gives  a  continuous  series  of  colors 
from  red  through  orange,  yellow, 
green,  blue,  to  violet,  called  a  con- 
tinuous spectrum. 

When  many  of  the  elements  (or 
their  compounds)  are  volatilized  at  a 
high  temperature,  as  in  the  heat  of 
the  Bunsen  flame,  colored  lights  are 
obtained,  each  metal  having  its  own 
characteristic  color.  These  lights  differ  from  white  light  in  that  they 
are  not  made  up  of  so  many  different  wave  lengths ;  hence  when  passed 
through  the  prism  the  spectrum  obtained  is  not  continuous,  but  merely 
shows  those  colors  of  which  the  light  is  composed. 

That  these  colors  may  be  made  as  distinct  and  as  sharply  separated 
as  possible,  the  light  should  shine  upon  the  prism  through  a  very  narrow 
slit  in  a  screen,  arranged  so  as  to  be  parallel  with  the  axis  of  the  prism. 
The  colors  will  then  be  a  series  of  narrow  lines,  each  an  image  of  the 
slit,  spread  out  parallel  with  each  other.  An  instrument,  the  essential 
parts  of  which  are  a  prism,  a  screen  provided  with  a  narrow  slit,  and 
lenses  for  focusing  the  light  upon  the  slit  and  for  viewing  the  spectrum, 
is  called  a  spectroscope,  or  spectrometer.  Fig.  84  represents  a  simple 
form  of  such  a  spectroscope,  the  slit  being  seen  at  the  end  of  the  tube  B. 
The  light  from  the  Bunsen  burner  passes  through  the  slit  in  the  tube  B, 
is  separated  into  its  different  wave  lengths  by  the  prism  at  the  center  of 
the  spectroscope,  and  is  seen  by  looking  through  the  tube  A.  The  tube  C 
contains  a  scale  illuminated  by  a  flame  at  the  end  of  the  tube ;  this  scale 
makes  it  possible  to  fix  the  position  of  any  lines  in  the  spectrum.  Make 
the  following  tests  with  the  spectroscope: 

[128] 


FIG.  83.  When  passed  through  a 
prism  a  ray  of  white  light  is  sepa- 
rated into  the  different  colors  which 
together  compose  the  white  light 


First  use  the  luminous  Bunsen  flame.  In  this  the  light  is  caused  by 
solid  particles  of  carbon,  and  the  spectrum  of  this  will  be  a  continuous 
band  when  viewed  through  A.  Next  examine  the  spectrum  of  one  or 
more  of  the  following  metals  :  potassium,  sodium,  calcium,  strontium, 
barium,  using  in  each  case  a  volatile  compound  of  the  metal,  preferably 
the  chloride  or  nitrate.  To  do  this  the  nonluminous  Bunsen  flame  must 
be  used.  First  clean  the  platinum  wire  until  it  will  impart  no  color  what- 
ever to  the  flame.  Then  dip  the  loop  at  the  end  of  the  wire  into  a 
concentrated  solution  of  sodium  chloride  and  bring  the  loop  into  the 


FIG.  84.  A  spectroscope 

lower  edge  of  that  part  of  the  Bunsen  flame  which  is  next  to  the  tube  B, 
as  shown  in  the  figure.  The  spectrum  will  be  found  to  consist  of  a  single 
band  of  yellow  light.  The  slit  at  the  end  of  the  tube  B  should  be  ad- 
justed by  the  attached  screw  until  the  band  is  narrow  and  sharp.  In  the 
same  way  examine  the  spectrum  produced  by  the  other  metals  mentioned 
above.  It  is,  of  course,  essential  that  the  wire  be  thoroughly  cleaned 
each  time;  otherwise  we  will  obtain  the  spectrum  of  two  or  more  of 
the  metals.  It  will  be  found  that  with  the  exception  of  sodium  the  other 
metals  mentioned  will  give  more  than  one  band  of  light,  (m)  Diagram 
the  results  obtained  for  each  of  the  metals. 


[129} 


EXERCISE  82 

THE  PROPERTIES  OF  BLEACHING  POWDER 

Apparatus.  2  small  beakers ;  stirring-rod ;  carbon  dioxide  generator 
(Fig.  53);  funnel;  test  tube. 

Materials.  25  g.  bleaching  powder  (see  that  the  bleaching  powder 
comes  from  air-tight  cans;  otherwise  the  powder  is  apt  to  be  worth- 
less) ;  10  g.  marble  and  hydrochloric  acid  for  generating  carbon  dioxide  ; 
filter  paper;  narrow  strips  of  different  samples  of  colored  calico. 

1.  The  active  agent  of  commercial  bleaching  powder  is  the 
compound  CaOCl2.  In  addition  to  this  compound  the  com- 
mercial powder  contains  a  certain  amount  of  calcium  hydroxide 
and  certain  other  impurities  present  in  the  lime  from  which  the 
bleaching  powder  was  made. 

When  sulfuric  acid  is  added  to  bleaching  powder  both  hydro- 
chloric acid  and  hypochlorous  acid  (HC1O)  are  formed.  These 
two  react  as  fast  as  generated,  however,  to  form  water  and 
chlorine.  The  equations  for  the  reactions  are  as  follows : 

CaOCl2  +  H2S04  — +  CaS04  +  HC1  +  HC1O 
HC1  +  HC1O  — +  H2O  +  C12 

A  very  weak  acid  like  carbonic  acid,  on  the  other  hand,  liber- 
ates only  hypochlorous  acid  as  follows : 

2CaOCl2  +  H2CO8  — >  CaCO3  +  CaCl2  -I-  2HC1O 

This  is  the  reaction  that  takes  place  when  bleaching  powder  is 
exposed  to  the  air,  the  carbonic  acid  being  formed  from  the 
moisture  and  carbon  dioxide  present  in  the  air.  Hypochlorous 
acid  is  a  good  bleaching  and  disinfecting  agent.  To  test  its 
bleaching  property,  proceed  as  follows : 

Pour  50  cc.  of  water  over  2  5  g.  of  bleaching  powder  in  a 
beaker.  Stir  the  mixture  for  two  or  three  minutes ;  then  filter. 
The  active  agent  of  the  powder,  CaOCl2,  is  now  present  in  the 
filtrate.  Now  pass  through  the  filtrate  a  slow  current  of  carbon 

[130] 


dioxide  as  long  as  a  precipitate  continues  to  form ;  then  filter. 

(a)  What  compounds  are  present  in  the  filtrate?    Immerse 
some  strips  of  colored  cloth  in  the  filtrate  for  a  few  minutes. 

(b)  Record  your  results. 

EXERCISE  83 

HARD  WATERS  AND  METHODS  FOR  SOFTENING  THEM 

Apparatus.  6o-cc.  bottle;  carbon  dioxide  generator  (Fig.  53);  fun- 
nel; 2  test  tubes;  small  beaker. 

Materials.  30  cc.  limewater  (R.S.)  ;  bit  of  soap  dissolved  in  water; 
10  to  i5g.  of  marble  and  hydrochloric  acid  for  generating  carbon 
dioxide;  i  g.  magnesium  sulfate  (R.S.)  ;  filter  paper;  i  g.  sodium  car- 
bonate dissolved  in  as  little  water  as  possible ;  i  g.  powdered  calcium 
sulfate. 

1.  Bubble  carbon  dioxide  into  250:.  of  limewater  diluted 
with  an  equal  volume  of  water.    Note  that  a  precipitate  forms 
which  gradually  dissolves  as  more  of  the  gas  is  passed  through. 
(a)  Account  for  the  results,  writing  the  equations  for  the  re- 
actions involved.   The  resulting  liquid  is  a  good  sample  of  a 
hard  water. 

Add  a  few  drops  of  a  soap  solution  to  5  cc.  of  the  hard  water 
and  (b)  note  the  results. 

Divide  the  remainder  of  the  hard  water  into  two  parts. 
Gradually  boil  the  one  part,  (c)  Note  and  account  for  the 
result.  To  the  other  part  add  a  few  drops  of  clear  limewater 
and  mix  intimately,  (d)  Again  note  and  account  for  the  result. 

2.  Shake  i  g.  of  calcium  sulfate  with  10  cc.  of  water  in  a 
test  tube  for  two  or  three  minutes;   filter,  and  add  to  the 
filtrate  2  or  3  drops  of  a  saturated  solution  of  sodium  carbon- 
ate,   (e)  Note  and  account  for  the  results. 

All  hard  waters  contain  more  or  less  calcium  acid  carbonate, 
calcium  sulfate,  calcium  chloride ;  also  the  corresponding  com- 
pounds of  magnesium.  (The  methods  used  for  removing  the 
calcium  compounds  likewise  serve  for  removing  the  magnesium 
compounds.)  (/)  How  could  such  waters  be  softened  on  a 

[131] 


large  scale?    (g)  Waters  softened  in  this  way  would  contain 
what  compounds  in  solution  ? 

3.  Test  some  hard  waters  from  wells  by  the  above  methods 
and  (h)  give  the  results  obtained. 

EXERCISE  84 
TESTING  THE  ACIDITY  OF  SOILS 

Apparatus.  Watch  glass;  ordinary  kitchen  knife  with  thin  blade  so 
that  it  will  readily  bend;  forceps;  (25o-cc.  flask;  tripod). 

Materials.  Cupful  of  soil  to  be  tested ;  strips  of  blue  and  red  litmus 
papers;  (i  g.  barium  chloride  mixed  with  o.i  g.  zinc  sulfide;  strip  of 
filter  paper  (about  6  cm.  x  i  cm.)  moistened  with  a  solution  of  lead  ace- 
tate (R.S.)). 

1.  Soils  under  cultivation  tend  to  become  acid,  and  most 
crops  do  not  thrive  in  acid  soil;  hence  the  custom  of  adding 
lime  to  soils.  The  following  is  a  simple  way  of  finding  out 
whether  or  not  a  given  soil  is  acid : 

Collect  about  one  cupful  of  an  average  sample  of  the  soil. 
By  means  of  a  clean  knife  blade  mix  the  soil  intimately.  Now 
take  one  strip  each  of  red  and  blue  litmus  paper  (use  your  for- 
ceps to  avoid  any  secretion  from  your  fingers  from  coming  in 
contact  with  the  papers) ,  moisten  the  papers  with  distilled  water 
and  press  them  side  by  side  against  the  concave  face  of  a  watch 
glass.  Pour  over  the  strips  2  or  3  tablespoonfuls  of  the  soil ; 
then  add  distilled  water,  a  little  at  a  time,  and  gently  mix  the 
water  through  the  soil  with  a  clean  knife  blade  (be  careful  not 
to  displace  the  red  and  blue  litmus  papers)  until  there  is 
formed  a  thick  paste  which  will  cling  to  the  watch  glass  when 
inverted.  The  glass  is  now  inverted  and  placed  on  the  desk 
so  that  the  strips  of  litmus  paper  are  visible.  After  fifteen  or 
twenty  minutes  note  if  any  change  in  the  color  of  the  litmus 
paper  has  taken  place,  (a)  Record  the  results  of  your  tests  of 
different  soils,  (b)  What  precautions  must  be  taken  in  order  to 
secure  reliable  results? 

[132] 


2.  A  method  of  testing  soils  largely  used  at  the  present  time  was  de- 
vised by  Professor  Truog  of  the  University  of  Wisconsin  and  is  known 
as  the  "  Truog  test."  Proceed  as  follows  : 

Collect  a  sample  of  soil  as  in  experiment  i  and  introduce  12  or  15  g. 
of  the  soil  into  a  25o-cc.  flask.  Next  add  to  this  i  g.  of  the  barium 
chloride — zinc  sulfide  mixture  (see  materials).  Pour  into  the  flask 
100  cc.  of  distilled  water  and  shake  the  flask 
very  gently.  Heat  the  mixture,  using  a  flame 
that  will  cause  the  liquid  to  boil  in  from  five 
to  seven  minutes.  Sometimes  the  liquid 
froths  just  when  it  begins  to  boil;  in  case  it 
begins  to  froth,  withdraw  the  heat  for  a 
moment  until  the  frothing  subsides.  After 
the  liquid  in  the  flask  has  boiled  for  just  one 
minute,  place  across  the  mouth  of  the  bottle, 
as  shown  in  Fig.  85,  a  strip  of  lead  acetate 
paper  and  hold  it  in  this  position  for  two 
minutes,  while  the  boiling  continues;  then 
remove  the  paper  and  withdraw  the  heat.  If 
the  soil  is  acid,  hydrogen  sulfide  will  be  lib- 
erated and  will  act  upon  the  lead  acetate, 
forming  black  lead  sulfide  (PbS).  The 
greater  the  acidity,  the  greater  the  amount 
of  lead  sulfide  formed,  and  hence  the  darker 

the  paper.  Since  the  test  is  very  delicate  it  is  essential  that  care  be  taken 
to  use  distilled  water  and  to  see  that  the  materials  are  pure  and  that  all 
apparatus  used  is  clean.  If  there  is  any  doubt  as  to  the  purity  of  the 
materials,  a  blank  test  should  be  made.  To  do  this  proceed  just  as 
directed  above,  except  that  the  sample  of  soil  is  omitted. 

EXERCISE  85 

ACTION  OF  HARD  WATERS  ON  SOAP  (OPTIONAL) 

Apparatus.  Two  2^o-cc.  bottles;  burette  or  graduated  cylinder; 
2  test  tubes. 

Materials,  i  g.  soap  dissolved  in  100  cc.  distilled  water ;  samples  of 
hard  water  (100  cc.);  samples  of  one  or  more  washing-powders. 

1.  The  determination  of  the  amount  of  soap  lost  by  using 
hard  water  for  washing.  Place  two  2  50-00.  bottles  on  the  desk. 
Into  the  first  pour  100  cc.  of  hard  water  (preferably  an  average 

[133] 


FIG.  85.    Testing  soils  by 
the  Truog  method 


sample  of  the  water  used  in  your  town  or  city)  and  into  the 
second  pour  100  cc.  of  distilled  (or  rain)  water.  Now  add  to 
each  the  soap  solution,  i  cc.  at  a  time,  and  shake  the  bottle 
vigorously  after  each  addition.  Continue  adding  the  soap  solu- 
tion until  the  lather  formed  on  shaking  the  mixture  persists 
for  five  minutes,  (a)  Compare  the  amounts  of  the  soap  solu- 
tion required  in  each  case  to  produce  a  permanent  lather.  The 
difference  in  the  amounts  represents  the  soap  consumed  by 
using  the  hard  water  in  place  of  soft  water. 

It  will  be  interesting  to  make  a  rough  approximation  of  the 
amount  of  hard  water  used  yearly  in  your  city  for  washing,  and 
then  to  determine  approximately  the  cost  of  the  soap  lost  in  one 
year  owing  to  the  use  of  hard  water. 

2.  The  analysis  of  washing-powders.  Devise  methods  for 
detecting  the  presence  of  the  following  substances,  if  present  in 
washing-powders:  (b)  sodium  carbonate;  (c)  borax  (Exer- 
cise 72);  (d)  mineral  matter,  such  as  sand.  Submit  your 
methods  to  your  instructor  for  criticism ;  then  test  one  or  more 
washing-powders  for  these  substances,  (e)  Record  your  results. 

EXERCISE  86 
MAGNESIUM  AND  ITS  COMPOUNDS 

Apparatus.  Porcelain  crucible  and  cover;  small  beaker;  pipestem 
triangle;  evaporating-dish ;  forceps. 

Materials.  Strips  of  magnesium  5  cm.  long ;  blue  and  red  litmus 
papers ;  2  to  3  g.  magnesium  carbonate ;  hydrochloric  acid. 

1.  Wind  a  strip  of  magnesium  wire  into  a  coil  and  place  it 
in  a  porcelain  crucible.  Put  the  cover  on  the  crucible  and  apply 
a  gentle  heat.  By  means  of  your  forceps  raise  the  cover  slightly 
from  time  to  time  so  as  to  admit  air.  Continue  until  the  mag- 
nesium is  entirely  burned,  leaving  a  white  powder,  (a)  What 
is  the  composition  of  the  powder?  Add  the  powder  to  25  cc. 
of  water.  Stir  the  mixture  and  test  it  with  litmus  paper. 
(b)  Account  for  the  results. 

[1341 


2.  Convert  2  or  3  g.  of  magnesium  carbonate  into  the  chloride. 
(c)  Describe  the  process.  Evaporate  the  solution  of  the  chlo- 
ride to  complete  dryness  in  an  evaporating-dish,  heating  the 
residue  gently  with  the  bare  flame.  When  it  is  cool  add  a  few 
drops  of  water,  stir,  and  test  with  litmus,  (d)  Account  for  the 
results,  (e)  Why  are  waters  containing  magnesium  chloride 
objectionable  for  use  in  steam  boilers?  (Consult  text.) 


EXERCISE  87 
ZINC  AND  ITS  COMPOUNDS 

Apparatus.    Blowpipe  and  charcoal  (Fig.  86);  3  test  tubes;  beaker. 
Materials.   0.5  g.   zinc;    sulfuric  acid;    sodium  hydroxide  solution; 
ammonium  sulfide  (R.S.);  i  g.  zinc  sulfate. 

1.  Place  a  bit  of  zinc  on  charcoal  and  heat  it  in  the  oxidizing 
flame  produced  by  a  blowpipe  (Fig.  86).    The  resulting  oxide 
is  deposited  as  a 

film  on  the  char- 
coal, (a)  Note  its 
color,  (b)  Is  its 
color  the  same 
when  hot  as  when 
cold?  (c)  For 
what  is  this  oxide 
used? 

2.  Dissolve  o.5g. 
of  zinc  sulfate  in 
10  cc.    of    water. 
Divide  this  solu- 
tion into  2  parts 
and      test      with 
the  following  re- 


FIG.  86.   Heating  a  metal  in  the  oxidizing  flame 


agents:  (i)  sodium  hydroxide  solution  (i  drop,  or  just  sufficient 
to  cause  a  precipitate) ;  zinc  hydroxide  precipitates  ((d)  write 
the  equation  for  its  formation),  but  the  precipitate  dissolves 

[1351 


again  if  an  excess  of  sodium  hydroxide  is  added;  (2)  am- 
monium sulfide;  zinc  sulfide  precipitates,  (e)  Write  the  equa- 
tion for  the  reaction  involved.  (/)  What  is  the  color  of  zinc 
sulfide? 

3.  Perform  experiment  i,  Exercise  88,  so  that  the  exercise 
may  be  concluded  at  your  next  laboratory  period. 


EXERCISE  88 

RUBBER 

Apparatus.  2  beakers  (loo-cc.  and  300-00.);  test  tube;  stirring-rod; 
glass  plate  for  covering  bottles. 

Materials.  1.5  g.  natural  rubber;  35  cc.  carbon  tetrachloride  ;  0.2  g. 
flowers  of  sulfur  or  finely  powdered  brimstone;  acetic  acid  (R.S.); 
filter  paper. 

1.  Put  about  i  g.  of  natural  rubber 
in  a  loo-cc.  beaker  and  pour  over  it 
25cc.  of  carbon  tetrachloride.   Place 
over  the  top  of  the  beaker  a  filter 
paper,  and  on  this  a  glass  plate  so 
as  to  prevent  the  liquid  from  evapo- 
rating.   Set  the  beaker  aside  until  the 
next  laboratory  period.    Also  intro- 
duce about  o.i  g.  of  natural  rubber 
into  a  test  tube  and  pour  over  it  suffi- 
cient carbon  tetrachloride  to  half  fill 
the  tube ;  then  cork  the  tube  and  set 
it   aside   until   the   next   laboratory 
period  (if  opportunity  is  afforded,  it 
is  well  to  stir  the  mixtures  once  or 
twice  before  the  next  period ) .  In  both 

cases  the  rubber  forms  a  colloid  gel  in  the  carbon  tetrachloride. 

2.  Stir  the  mixture  in  the  beaker  thoroughly;  then  add  to 
it  about  0.2  g.  flowers  of  sulfur  and  stir  the  sulfur  until  it  is 
uniformly  distributed  through  the  mass.    Pour  into  a  large 

[136] 


FIG.  87.  Evaporating  a  liquid 
by  placing  the  beaker  contain- 
ing the  liquid  inside  a  larger 
beaker  partially  filled  with 
boiling  water 


beaker  sufficient  water  so  that  the  smaller  beaker  containing 
the  rubber  will  float  on  the  water,  as  shown  in  Fig.  87.  Heat 
the  water  to  boiling  and  maintain  the  boiling  for  half  an  hour 
or  more,  until  the  carbon  tetrachloride  in  the  small  beaker  is 
all  evaporated  and  the  rubber  forms  a  film  closely  adhering  to 
the  sides  of  the  beaker.  Then  remove  the  small  beaker,  and 
by  means  of  a  knife  blade  peel  off  the  adhering  rubber. 

(a)  Compare   the  resulting  rubber   in  properties   with  the 
natural  rubber. 

3.  Shake  the  test  tube  containing  the  colloid  rubber  vigor- 
ously so  as  to  form  a  uniform  mass;  then  add  2  or  3  cc.  of 
acetic  acid  and  again  shake  the  tube.  Set  the  tube  aside  for  a 
few  minutes,  (b)  Does  the  colloid  (rubber)  settle?  (c)  What 
is  the  action  of  the  acetic  acid?  (Consult  chapter  on  colloids 
in  text.) 

EXERCISE  89 

ALUMINIUM  AND  ITS  COMPOUNDS 

Apparatus.   2  test  tubes;  blowpipe  and  charcoal;  2  beakers;  funnel. 

Materials.  2  g.  aluminium  turnings ;  sodium  hydroxide  solution ; 
wooden  splint ;  hydrochloric  acid  ;  ammonium  hydroxide  ;  filter  paper  ; 
aluminium  sulfate  solution  (R.  S.) ;  2  or  3  drops  of  cobalt  nitrate  solu- 
tion (R.  S.) ;  i  g.  sodium  carbonate  in  5  cc.  water;  (aluminium  sulfate  and 
potassium  sulfate  sufficient  to  make  20  g.  of  crystals  of  potassium  alum). 

1.  (a)  Note  the  physical  properties  of  aluminium.    Introduce 
about  0.5  g.  of  aluminium  into  a  test  tube  and  pour  over  it  2  cc. 
of    sodium   hydroxide    solution.     Heat    the    mixture    gently. 

(b)  What  evidence  have  you  that  the  aluminium  is  acted  upon 
by  the  hydroxide?    (c)  Close  the  mouth  of  the  tube  for  a  few 
seconds  with  your  thumb ;  then  remove  it  and  at  once  test  the 
gas  with  a  lighted  splint,    (d)  What  is  the  gas?    (e)  What 
would  be  the  effect  of  washing  aluminium  cooking  vessels  with 
lye  or  strong  soap  ? 

2.  Introduce  about  i  g.  of  aluminium  into  a  test  tube,  add 
5  cc.  of  water  and  then  hydrochloric  acid,  a  drop  at  a  time, 

[137] 


sufficient  to  dissolve  the  metal.  Filter  (if  necessary)  and  dilute 
the  solution  to  about  50  cc.  To  this  add  ammonium  hydroxide 
until  the  solution  reacts  alkaline.  .(/)  Note  the  results  and 
write  the  equation  for  the  reactions  involved. 

3.  Filter  off  some  of  the  aluminium  hydroxide  prepared  in 
i  and  heat  it  on  charcoal  in  the  oxidizing  flame  of  the  blowpipe 
(Fig.  86).   The  aluminium  hydroxide  is  decomposed  into  alu- 
minium oxide  and  water,    (g)  Write  the  equation  for  the  re- 
action.  Now  moisten  the  residue  on  the  charcoal  with  one  or 
two  drops  of  a  solution  of  cobalt  nitrate  and  again  heat  strongly 
in  the  blowpipe  flame,    (h)  Note  the  change  in  color.   Ad- 
vantage is  sometimes  taken  of  this  property  in  testing  for  the 
presence  of  aluminium. 

4.  Add  a  solution  of  sodium  carbonate  to  any  aluminium  salt 
such    as   aluminium   sulfate.    Note    that    a    gas    is    evolved. 
(i)  Devise  a  method  for  determining  whether  or  not  this  gas 
is  carbon  dioxide  and  make  the  test.    (;)  Advantage  is  taken 
of  this  reaction  in  the  preparation  of  what  culinary  product  ? 

5.  Calculate  the  weights  of  aluminium  sulfate  (the  crystals  of  alu- 
minium sulfate  have  the  formula  A12(S04)3  •  16  H20)  and  of  potassium 
sulfate  required  to  prepare  20  g.  of  crystals  of  potassium  alum;  then  dis- 
solve these  amounts  of  the  two  compounds  separately  in  as  little  water 
as  possible,  mix  the  two  solutions  thoroughly,  and  set  the  resulting  solu- 
tion aside  for  a  few  days  to  crystallize.    If  a  string  is  suspended  in  the 
liquid,  the  crystals  will  deposit  on  it.    These  may  then  be  withdrawn 
and  their  properties  studied,    (k)  Note  the  shape  of  the  crystals. 

EXERCISE  90 

A  STUDY  OF  THE  USE  OF  ALUMINIUM  SULFATE  IN 
THE  PURIFICATION  OF  WATER 

Apparatus.    Three  25o-cc.  wide-mouthed  bottles ;  graduated  tube. 
Materials.     Aluminium  sulfate  solution  (R. S.)  ;  limewater  (R.S.). 

1.  Label  three  25o-cc.  wide-mouthed  bottles  A,  B,  and  C 
respectively.  Nearly  fill  A  and  B  with  muddy  water,  pouring 
a  like  volume  of  distilled  water  into  C.  Add  5  drops  of  alu- 

[  138] 


minium  sulfate  solution  to  A  and  C  respectively.  Mix  the  con- 
tents of  each  bottle  thoroughly.  Now  to  each  of  the  bottles  A 
and  C  add  10  cc.  of  limewater.  Set  the  bottles  aside  and 
examine  at  the  beginning  of  the  next  laboratory  period. 
(a)  Record  the  results  and  explain  the  reactions.  What  is  the 
use  of  bottle  B  ? 

EXERCISE  91 

REACTIONS  OF  BAKING-POWDERS 

Apparatus.    250-00.  flask  to  fit;  stirring-rod;  mortar  and  pestle. 
Materials.   4  g.  sodium  bicarbonate  ;  limewater  (R.  S.) ;  alum ;  cream 
of  tartar. 

1.  By  referring  to  the  equations  for  the  reactions  of  different 
classes  of  baking-powders  as  given  in  your  text,  calculate  and 
(a)   record  the  weight  of  ammonium  alum   (NH4A1(SO4)2  • 
i2H0O)  necessary  to  react  with  2  g.  of  sodium  bicarbonate; 
then  grind  together  these  weights  of  alum  and  bicarbonate  so 
as  to  mix  them  thoroughly.    Put  the  mixture  into  a  250-0:. 
flask  and  cover  the  mixture  with  water.    Gently  rotate  the 
flask  so  as  to  form  a  uniform  mixture,  at  the  same  time  heating 
the  mixture  slightly ;  then  set  the  flask  aside  for  five  minutes. 
Now  test  for  the  presence  of  carbon  dioxide  in  the  air  in  the 
flask  (Fig.  54).    (b)  Note  and  explain  the  results. 

2.  Repeat   i,   substituting  cream  of  tartar   for   the  alum. 
(c)  Record  your  results. 

3.  (d)  What  compounds  remain  in  food  as  a  result  of  the  use 
of  an  alum  baking-powder;   (e)  of  a  cream-of -tartar  baking- 
powder  ? 


139] 


EXERCISE  92 
ANALYSIS  OF  BAKING-POWDERS  (OPTIONAL) 

Apparatus.  2  small  beakers ;  stirring-rod ;  funnel ;  evaporating-dish ; 
5  test  tubes. 

Materials.  10  g.  each  of  various  kinds  of  baking-powders ;  iodine 
solution  (R.S.);  filter  paper;  barium  chloride  solution  (R.S.);  hydro- 
chloric acid;  sulfuric  acid;  nitric  acid;  ammonium  molybdate  solution 
(R.S.)  ;  sodium  hydroxide. 

1.  Introduce  10  g.  of  a  baking-powder  into  a  beaker  and  pour 
over  it  50  cc.  of  water.    Stir  the  mixture  thoroughly  until  no 
more  gas  is  evolved,  then  filter  it,  and  test  the  residue  and  the 
filtrate  for  the  various  ingredients,  as  explained  below. 

2.  Starch,    (a)  Will  any  starch  present  be  in  the  residue  or 
in  the  filtrate?    Make  appropriate  test,    (b)  Record  the  test 
used  and  the  results  obtained. 

3.  Sulfates.  (c)  All  alums  are  salts  of  what  acid?  (d)  Make 
appropriate  test  for  these  salts,  giving  the  method  and  the 
results. 

4.  Tartrates.    Pour  5  cc.  of  the  filtrate  into  an  evaporating- 
dish,  add  5  drops  of  sulfuric  acid,  and  evaporate  to  dryness. 
Finally,  heat  the  dish  gently  with  a  bare  flame.   The  presence 
of  a  tartrate  is  indicated  by  an  odor  similar  to  that  of  burning 
sugar,    (e)  Record  the  results  of  your  test  of  different  powders. 

5.  Ammonium  salts.    Ammonium  alum  is  sometimes  used  in 
baking-powders.   To  detect  this,  pour  5  cc.  of  the  filtrate  into 
a  test  tube,  add  an  equal  volume  of  sodium  hydroxide  solution, 
and  heat  gently.   If  ammonium  salts  are  present,  ammonia  will 
be  evolved  (Exercise  78). 

6.  Phosphates.    If  calcium  or  sodium  phosphate  is  present  in 
the  baking-powder,  the  filtrate  will  contain  the  acid  phosphate 
of  the  metal.   To  detect  phosphates,  treat  5  cc.  of  the  filtrate 
with  a  few  drops  of  nitric  acid,  heat  nearly  to  boiling,  and  add  a 

[140] 


few  drops  of  the  mixture  to  5  cc.  of  ammonium  molybdate  solu- 
tion (compare  Exercise  67).  (/)  Record  the  results  obtained. 

7.  Aluminium,  calcium,  sodium,  and  potassium.  A  baking- 
powder  containing  sulfates  always  contains  aluminium,  while 
one  containing  phosphates  always  contains  calcium  or  sodium ; 
cream-of-tartar  baking-powders,  on  the  other  hand,  always 
contain  potassium.  The  presence  or  absence  of  these  metals 
may  be  inferred  from  the  tests  made  for  sulfates,  phosphates, 
and  cream  of  tartar. 

EXERCISE  93 

THE  USE  OF  MORDANTS  IN  DYEING  (OPTIONAL) 

Apparatus.    200-cc.  beaker;  stirring-rod;  large  beaker. 

Materials.  2  strips  (2  cm.  x  6  cm.)  of  white  woolen  cloth  (nun's 
veiling  serves  well) ;  6  strips  of  white  cotton  cloth ;  i  g.  sodium  car- 
bonate in  50  cc.  water ;  0.5  g.  tannic  acid  in  50  cc.  water  ;  0.2  g.  tartar- 
emetic  in  50  cc.  water;  o.ig.  of  any  of  the  following  dyes  in  1500:. 
water  (different  students  should  select  different  dyes  and  compare  re- 
sults):  fuchsine,  methyl  violet,  gallein,  malachite  green,  Congo  red;  (a 
solution  containing  i  g.  sodium  carbonate,  5  g.  Glauber's  salt,  and  o.i  g. 
Congo  red  in  50  cc.  water). 

1.  Most  dyes  will  dye  animal  fibers  (wool,  silk)  directly, 
but  will  dye  vegetable  fibers  (cotton,  linen)  fast  only  when 
mordants  are  used. 

Place  the  strips  of  cotton  cloth  in  a  beaker  and  cover  them 
with  the  sodium  carbonate  solution  ( i  g.  in  50  cc.  of  water) 
and  boil  the  liquid  for  five  minutes.  Remove  the  strips  and 
thoroughly  rinse  them  with  water.  This  treatment  serves  to 
remove  all  foreign  matter  from  the  cloth. 

Now  completely  immerse  two  strips  of  the  cotton  cloth  in 
the  tannic-acid  solution,  heat  until  it  is  fairly  warm  to  the  hand 
(50°  or  60°),  and  maintain  the  temperature  for  ten  minutes. 
Remove  the  cloth  from  the  solution  and  squeeze  out  the  liquid, 
but  do  not  rinse  the  cloth;  then  immerse  the  strips  for  one 
minute  in  the  slightly  warmed  solution  of  tartar-emetic.  The 

[141] 


tannic  acid  in  the  cloth  reacts  with  the  tartar-emetic,  forming 
a  salt  (known  as  antimonyl  tannate)  which  becomes  incor- 
porated in  the  meshes  of  the  fiber  and  serves  as  a  mordant. 
Remove  the  cloth  from  the  solution  and  rinse  it. 

Divide  the  solution  of  the  dye  chosen  into  three  equal  por- 
tions. Heat  one  portion  in  a  small  beaker  just  to  boiling,  im- 
merse a  strip  of  woolen  cloth,  and  continue  the  heating  for 
about  five  minutes.  Be  sure  to  keep  the  cloth  entirely  im- 
mersed in  the  dye,  using  the  glass  stirring-rod  for  this  purpose. 
Now  remove  the  cloth  and  rinse  it  thoroughly.  To  test  whether 
the  cloth  is  dyed  fast,  rinse  it  thoroughly  and  then  wash  it  in  a 
beaker  of  water  and  note  whether  the  water  becomes  colored. 
In  a  similar  way  dye  a  strip  of  unmordanted  cotton  cloth  in  the 
second  portion  of  the  dye  and  a  strip  of  mordanted  cloth  in 
the  third  portion,  and  determine  in  each  case  whether  the  cloth 
is  dyed  fast,  (a)  Finally,  dry  the  three  strips  of  cloth  and  in- 
sert them  in  your  notebook  and  record  the  results  of  your 
experiments. 

2.  Some  dyes  (known  as  substantive  dyes)  have  the  property  of 
dyeing  cotton  fast  without  the  use  of  mordants.  The  ordinary  dyes 
sold  by  druggists  belong  to  this  class.  Congo  red  is  a  typical  dye  of 
this  class. 

Heat  the  solution  of  Congo  red,  prepared  as  directed  under  "Mate- 
rials," to  boiling,  immerse  in  it  a  strip  of  wet  unmordanted  cotton,  and 
continue  the  boiling  for  five  minutes.  Remove  the  cloth,  rinse,  and 
test  to  see  whether  the  cloth  is  dyed  fast.  Dry  the  strip  and  insert 
it  in  your  notebook.  (The  sodium  carbonate  and  the  sodium  sulfate 
assist  in  the  process,  but  do  not  act  as  mordants.) 

(Save  the  remaining  strips  of  cotton  for  use  in  Exercise  94.) 


[142] 


EXERCISE  94 

A  STUDY  OF  LAKES;  ALSO  THE  EFFECT  OF  USING  DIF- 
FERENT MORDANTS  WITH  THE  SAME  DYE  (OPTIONAL) 

Apparatus.  Two  250-00.  wide-mouthed  bottles;  evaporating-dish ; 
2  small  beakers;  stirring-rods. 

Materials.  5  cc.  of  a  2o-per-cent  solution  of  alizarin  paste;  o.i  g. 
of  gallein  dissolved  in  50  cc.  water ;  solutions  of  aluminium  sulfate  and 
ferric  sulfate  (R. S.);  ammonium  hydroxide;  3  strips  of  cotton  cloth 
prepared  in  Exercise  93. 

1.  Formation  of  lakes.   Label  two  wide-mouthed  bottles  A 
and  B  respectively.    Shake  the  alizarin  paste  so  as  to  form  a 
uniform  mixture;  then  introduce  about  10  drops  of  the  paste 
into  each  of  the  bottles.    Next  add  to  each  of  the  bottles  2  cc. 
of  ammonium  hydroxide  and  then  200  cc.  of  water,  and  mix  the 
contents  thoroughly.    Now  add  10  cc.  of  aluminium  sulfate 
solution  to  bottle  A  and  10  cc.  of  ferric  sulfate  to  bottle  B, 
intimately  mix  the  solutions,  and  set  the  bottles  aside;  (a)  note 
the  appearance  of  the  contents  at  the  end  of  the  laboratory 
period,  also  at  the  beginning  of  the  next  laboratory  period. 
(b)  What  is  the  function  of  each  of  the  materials  used? 

2.  Mordanting  strips  of  cloth  with  aluminium  hydroxide  or 
ferric  hydroxide.    Pour  about  20  cc.  of  aluminium  sulfate  solu- 
tion or  ferric  sulfate  solution  into  a  small  beaker  and  heat  to 
boiling  (different  students  should  use  different  solutions  and 
compare  results).    Completely  immerse  in  this  solution  two  of 
the  strips  of  cotton  cloth  prepared  in  Exercise  93,  and  continue 
the  heating  for  from  two  to  three  minutes.    Remove  the  cloth, 
squeeze  it  between  the  fingers  to  remove  the  excess  of  the 
solution,  and  immerse  it  in  20  cc.  of  water  containing  from  i 
to  2  cc.  of  ammonium  hydroxide.    Warm  the  liquid  slightly 
for  two  minutes,  then  remove  the  cloth  and  rinse  it  twice  in 
water,    (c)  What  compound  is  now  incorporated  in  the  cloth? 

f  1431 


3.  Dyeing  a  strip  of  mordanted  cloth  with  alizarin.   Pour 
i  cc.  of  alizarin  paste  into  a  small  beaker,  add  20  cc.  of  water, 
stir,  and  heat  to  boiling.    Completely  immerse  one  of  the  strips 
of  mordanted  cloth  prepared  in  2  and  continue  the  heating  and 
stirring  for  five  minutes;   then 'remove  the  cloth  and  rinse 
thoroughly.   Dry  the  dyed  strip  and  insert  it  in  your  notebook. 

4.  Dyeing  a  strip  of  mordanted  cloth  with  gallein.     Heat 
the  solution  of  gallein  just  to  boiling ;  then  immerse  one  of  the 
mordanted  strips  of  cotton  in  the  dye.    Keep  the  strip  com- 
pletely immersed  in  the  dye,  and  continue  the  heating  for  five 
minutes;  then  remove  the  cloth  and  rinse  it.   Dry  the  dyed 
strip  and  insert  it  in  your  notebook. 

EXERCISE  95 
THE  DETECTION  OF  DYES  IN  FOODS  (OPTIONAL) 

Apparatus.  Small  beaker;  stirring-rod. 

Materials.  Samples  of  colored  pop  and  orangeade ;  strips  of  woolen 
cloth  (nun's  veiling  gives  good  results) ;  hydrochloric  acid. 

1.  Select  different  samples  of  colored  pop.    Pour  50  cc.  of 
each  into  a  beaker,  add  2  or  3  drops  of  hydrochloric  acid, 
and  heat  to  boiling ;  then  introduce  a  strip  of  woolen  cloth  and 
continue  the  heating  for  -five  minutes.    Be  careful  to  keep  the 
cloth  completely  immersed  in  the  liquid,    (a)    Remove  the 
cloth,  rinse,  and  note  the  color. 

2.  Samples  of  colored  candies  may  be  tested  by  first  dissolv- 
ing the  candy  in  water  and  then  testing  the  solution  for  dyes 
by  using  strips  of  woolen  cloth,  as  in  the  above  case. 

Tomato  catchup  is  sometimes  colored,  although  the  practice  is  for- 
bidden by  Federal  law.  To  test  a  catchup  for  dyes,  heat  a  portion  of  the 
catchup  diluted  with  water  and  immerse  a  strip  of  woolen  cloth  in  the 
hot  mixture  for  five  minutes.  Remove  the  cloth  and  rinse  thoroughly. 
If  artificial  dyes  are  present,  the  cloth  will  be  deeply  colored;  otherwise 
it  will  have  only  a  slight  brownish  tinge  produced  by  the  natural  coloring- 
matter  of  the  tomato. 

[144] 


EXERCISE  96 
CLAY;  PORTLAND  CEMENT;  MORTAR 

Apparatus.  Evaporating-dish  ;  2  beakers ;  a  piece  of  wooden  shingle  or 
board  about  as  large  as  the  palm  of  your  hand ;  ordinary  kitchen  knife. 

Materials.  Small  samples  (half  a  cupful)  of  samples  of  different 
clays;  30 g.  Portland  cement;  15  g.  sand;  15 g.  lime  (this  should  be 
taken  from  a  lump  of  lime  freshly  made);  2  small  pasteboard  boxes 
(pill  boxes  will  do). 

1.  Add  water,  a  little  at  a  time,  to  different  samples  of  clay 
and  work  it  intimately  through  the  clay  until  a  thick,  plastic 
mass  is  obtained.  Devise  a  rough  method  for  testing  the  relative 
plasticity  of  different  clays  in  a  general  way  and  apply  it  to 
the  clays  at  hand,  arranging  them  in  the  order  of  their  plas- 
ticity,   (a)  Describe  the  method  you  used. 

2.  To  25  or  30  g.  of  Portland  cement  add  water,  a  little  at  a 
time,  and  stir  it  through  the  mass  until  a  thick  paste  is  obtained. 
Spread  a  portion  of  this  on  a  shingle  and  set  it  aside  until  the 
next  laboratory  period.   With  the  remainder  fill  a  small  paper 
pill  box.   Place  the  box  in  a  beaker  half  full  of  water.   Set 
the  beaker  and  contents  aside  until  the  next  period,    (b)  Ex- 
amine the  cement  placed  on  the  board,  also  that  in  the  beaker, 
and  note  whether  the  samples  have  hardened,  or  "set." 

3.  Repeat  experiment  2,  using,  in  place  of  the  cement,  some 
mortar  made  as  follows:  Add  to  15  g.  of  lime  10  cc.  of  water. 
The  lime  will  "slake"  and  form  a  thin  paste  with  the  excess  of 
water.   Now  work  through  this,  a  little  at  a  time,  15  g.  of  fine 
sand,  adding  more  water  if  necessary,  to  keep  the  mortar  from 
becoming  too  thick.    Place  some  of  this  on  a  shingle  and  with 
the  rest  fill  a  paper  pill  box;  then  proceed  as  in  experiment  2. 
(c)  Describe  the  results  obtained,  noting  in  particular  any 
difference  in  the  properties  of  Portland  cement  and  mortar. 

[145] 


EXERCISE  97 
A  STUDY  OF  IRON  AND  ITS  COMPOUNDS 

Apparatus.  Forceps;  2  beakers;  watch  glass;  funnel;  (flask  (25o-cc.); 
8  test  tubes). 

Materials.  Watch  spring  from  10  to  15  cm.  in  length  (broken  watch 
springs  can  be  obtained  from  any  jeweler)  ;  5  g.  small  tacks  or  fine 
iron  wire;  0,5  g.  powdered  iron;  filter  paper;  sulfuric  acid;  (nitric  acid; 
hydrochloric  acid ;  ammonium  hydroxide ;  potassium  ferrocyanide 
(R.S.);  potassium  sulfocyanate  (R.S.)). 

1.  The  tempering  of  steel.   Heat  a  piece  of  watch  spring 
(from  10  to  15  cm.  in  length)  to  a  white  heat  in  a  Bunsen  flame. 
Let  it  cool  slowly,  and  when  cold  bend  it  to  determine  if  it  is 
brittle.   Again  heat  to  a  white  heat  and  at  once  plunge  into  a 
beaker  of  cold  water.   When  cool,  bend  the  piece  as  before. 
Reheat  the  piece,  allow  it  to  cool  slowly,  and  again  examine  it. 
(a)  Record  the  results  of  the  experiments. 

2.  Preparation  of  ferrous  sulfate  (copperas).    Place  5  g.  of 
fine  iron  wire  or  small  tacks  in  a  beaker  and  pour  over  it  15  cc. 
of  water.   Now  add  (care)  4  cc.  of  concentrated  sulfuric  acid 
and  heat  very  gently  (hood)  until  a  vigorous  evolution  of  gas 
takes  place ;  then  cover  the  beaker  with  a  watch  glass  and  set 
it  aside  in  the  hood  until  near  the  end  of  the  laboratory  period. 
Then  add  10  cc.  of  water,  and  heat  slowly  until  the  liquid  boils, 
stirring  the  mixture  constantly.    Filter  off  any  undissolved 
solids,  collecting  the  nitrate  in  a  beaker.    Set  the  uncovered 
beaker  containing  the  filtrate  in  your  desk  until   the  next 
laboratory  period,    (b)  Record  the  properties  of  the  crystals. 
(c)  Write  the  equation  for  their  formation  from  iron. 

3.  Preparation  of  ferrous  and  ferric  salts  and  their  reactions.    Place 
about  0.5  g.  of  iron  powder  in  a  small  flask  and  pour  over  it  5  cc. 
of  water  and  then  2  cc.  of  hydrochloric  acid.   Mix  the  contents  of  the 
flask,  heat  the  flask  gently,  and  set  it  aside  in  the  hood  for  about  five 
minutes.   The  iron  dissolves  in  the  hydrochloric  acid,  forming  ferrous 

r  1461 


chloride.  Now  add  50  cc.  of  water  to  the  flask,  mix  well,  and  filter  off 
the  undissolved  iron.  Nearly  fill  a  test  tube  with  the  filtrate,  add  to  it 
4  or  5  drops  of  hydrochloric  acid  and  a  small  piece  of  iron  wire  or  i  or 
2  tacks,  and  loosely  cork  the  tube.  Mark  this  "Solution  A."  It  consists 
of  a  solution  of  ferrous  chloride.  The  iron  wire  and  the  hydrochloric 
acid  generate  a  little  hydrogen,  which  prevents  the  oxidation  of  the 
ferrous  chloride  to  ferric  chloride  by  the  oxygen  of  the  air.  Mark  the 
remainder  of  the  filtrate  "Solution  B." 

To  Solution  B  add  i  cc.  of  hydrochloric  acid  and  heat  it  nearly  to 
boiling;  then  withdraw  the  flame  and  add  nitric  acid,  drop  by  drop,  with 
constant  stirring,  until  the  solution,  which  is  at  first  dark  brown  in  color, 
becomes  light  in  color  (about  2  cc.  of  nitric  acid  will  be  required) 
The  ferrous  chloride  in  the  solution  is  changed  to  ferric  chloride  by  the 
oxygen  furnished  by  the  nitric  acid,  thus : 


2  FeCl2  +  2  HC1  +  0 


2  FeCl 


H20 


Now  compare  the  action  of  the  following  reagents  upon  Solutions  A  and 
B  (add  2  or  3  drops  of  the  reagents  to  3  cc.  of  the  solutions  in  separate 
test  tubes) :  ammonium  hydroxide,  potassium  ferrocyanide,  potassium 
sulfocyanate  (KCNS).  Tabulate  your  results  as  follows  : 


FERROUS  CHLORIDE 
(Solution  A) 

FERRIC  CHLORIDE 
(Solution  B) 

Ammonium  hydroxide  

Potassium  ferrocyanide      .... 
Potassium  sulfocyanate      .... 

EXERCISE  98 
THE  REMOVAL  OF  STAINS 

Apparatus.    4  test  tubes;  2  beakers;  evaporating-dish. 

Materials.  0.2  g.  tannic  acid  dissolved  in  10  cc.  water ;  ferric  sulfate 
(R.  S.) ;  2  pieces  of  white  cloth  10  cm.  square ;  i  g.  oxalic  acid  dissolved 
in  50  cc.  water  ;  hydrochloric  acid  ;  2  strips  of  black  cloth  5  cm.  square ; 
nitric  acid ;  ammonium  hydroxide ;  a  few  drops  each  of  cottonseed  oil 
and  sirup  (molasses)  ;  25  cc.  carbon  tetrachloride  ;  blotting-paper;  strips 
of  cloth  stained  witH  coffee  and  fruit  juice ;  acetic  acid  (R.  S.) ;  20  g. 
bleaching-powder ;  hot  water. 

1.  Stain  two  strips  of  white  cloth  by  dipping  them  into  a  solu- 
tion of  ferric  sulfate  until  thoroughly  saturated  and  then  into 

[147] 


a  solution  of  tannic  acid  (or  they  may  be  stained  directly  with 
black  ink).  Wash  one  of  the  strips  repeatedly  with  boiling 
water.  Leave  the  other  exposed  to  the  air  until  dry ;  then  try 
the  effect  of  hot  water  upon  the  stain.  If  the  stain  is  not  re- 
moved, wash  with  a  dilute  solution  of  oxalic  acid  and  finally 
with  hot  water,  (a)  Record  your  results. 

2.  Place  2  or  3  drops  of  dilute  hydrochloric  acid  upon  a  strip 
of  black  cloth,    (b)  What  color  is  produced?   Wash  the  spots 
with  10  cc.  of  water  containing  4  or  5  drops  of  ammonium 
hydroxide,    (c)  Do  the  spots  disappear?    Repeat,  using  nitric 
acid  in  place  of  hydrochloric  acid,    (d)  Account  for  the  results. 

3.  Place  in  separate  test  tubes  4  or  5  drops  of  a  sirup  and  a 
like  amount  of  fat,  such  as  cottonseed  oil.   Test  the  solubilities 
of  each  in  water  and  in  carbon  tetrachloride.    (e)  Suggest  a 
method  for  removing  stains  made  by  sirups  and  one  for  those 
made  by  fats.    Stain  some  strips  of  cloth  with  sirup  and  with 
an  oil  and  test  your  methods  for  removing  these  stains.    (In 
applying  a  solvent  it  is  convenient  to  place  the  stained  portion 
of  the  cloth  over  a  piece  of  blotting-paper.   A  small  bit  of 
sponge  or  cloth,  saturated  with  the  solvent,  is  then  rubbed 
about  the  stained  portion,  gradually  nearing  the  stain  itself, 
which  is  finally  thoroughly  rubbed.) 

Benzine  (or  gasoline)  may  be  used  in  place  of  the  carbon 
tetrachloride.  //  benzine  is  used,  however,  it  must  be  remem- 
bered that  it  is  very  inflammable.  Never  use  it  in  the  vicinity 
of  a  flame. 

4.  Stain  some  strips  of  cloth  with  coffee  and  with  fruit  juices. 
(/)  Are  the  stains  removed  by  washing  with  boiling  water? 
If  the  stain  cannot  be  removed  in  this  way,  wash  the  stained 
portion  of  the  cloth  in  bleaching-powder  to  which  has  been 
added  some  water  and  3  or  4  drops  of  acetic  acid. 


148 


EXERCISE  99 
A  STUDY  OF  COPPER  AND  ITS  COMPOUNDS 

Apparatus.  3  test  tubes  ;  beaker ;  (porcelain  crucible  ;  pipestem  tri- 
angle; balance). 

Materials.  Nail ;  copper  sulfate  solution  (R.  S.) ;  sodium  hydroxide  ; 
ammonium  hydroxide  ;  10  cm.  copper  wire ;  hydrochloric  acid  ;  (3  g. 
copper  sulfate  crystals). 

1.  (a)  Recall  the  action  of  nitric  acid  and  of  sulfuric  acid  on 
copper  (Exercises  37,  43) ;  also  the  action  of  sulfur  on  copper 
(Exercise  40).    Suspend  an  iron  nail  so  that  half  its  length  is 
immersed  in  a  solution  of  copper  sulfate  for  five  minutes. 
(b)  Account  for  the  result. 

2.  To  2  cc.  of  a  cold  solution  of  copper  sulfate  in  a  test  tube 
add  one  half  its  volume  of  sodium  hydroxide  solution.    Copper 
hydroxide,  Cu(OH)0,  is  precipitated.   Now  heat  to  boiling. 
The  hydroxide  is  decomposed  into  water  and  cupric  oxide 
(black),    (c)  Write  the  equations  for  the  reactions. 

3.  Add  i  drop  of  ammonium  hydroxide  to  a  dilute  solution 
of  copper  sulfate ;  now  continue  to  add  the  ammonium  hydrox- 
ide, drop  by  drop,  until  the  precipitate  which  is  at  first  formed 
is  dissolved,    (d)  How  does  the  color  of  this  solution  compare 
with  that  of  the  original  solution?    This  reaction  is  charac- 
teristic of  copper  compounds. 

4.  (e)  Recall  the  formation  of  cuprous  oxide  (Exercise  55). 

5.  Moisten  the  end  of  a  copper  wire  with  hydrochloric  acid 
and  hold  it  in  the  edge  of  a  Bunsen  flame.    (/)  What  is  the 
color  of  the  flame?    (g)  What  is  the  function  of  the  acid? 

6.  The  determination  of  the  percentage  of  water  of  hydration  in 
copper  sulfate  (quantitative).    Accurately  weigh   (or  counterpoise)   a 
porcelain  crucible  and  cover.    Record  all  weights  in  the  table  below. 
Place  2  or  3  g.  of  crystals  (no  larger  than  a  pea)  of  copper  sulfate  in 
the  crucible  and  again  accurately  weigh.    Place  the  covered  crucible  on 
a  pipestem  triangle  and  heat  it  with  a  gentle  flame  until  the  crystals  lose 

[149] 


their  color.  This  will  require  from  twenty  to  thirty  minutes.  The  tip 
of  the  flame  should  not  quite  touch  the  crucible.  The  product  is  an- 
hydrous copper  sulfate.  When  the  crucible  is  cool,  reweigh.  From  your 
results  calculate  the  percentage  of  water  of  crystallization  in  the  crys- 
tals. Fill  in  the  blank  spaces  in  the  table  below : 

Weight  of  crucible g. 

Weight  of  crucible  plus  copper  sulfate  crystals  ...  g. 

Weight  of  copper  sulfate  (calculate) g. 

Weight  of  crucible  plus  anhydrous  copper  sulfate     .      .  g. 

Weight  of  anhydrous  copper  sulfate g. 

Weight  of  water  of  hydration  (calculate)     ....  g. 

Percentage  of  water  of  hydration  in  crystals     ...  % 

Theoretical  percentage  of  water  of  hydration    ...  % 

Average  of  results  obtained  by  members  of  the  class  .  % 

EXERCISE  100 
A  STUDY  OF  MERCURY  AND  ITS  COMPOUNDS 

Apparatus.    loo-cc.  beaker;  2  test  tubes. 

Materials.  Globule  of  mercury  (size  of  a  grain  of  wheat) ;  nitric  add ; 
copper  penny  ;  0.5  g.  mercuric  oxide ;  3  cc.  solution  of  mercurous  nitrate 
(R.  S.);  hydrochloric  acid. 

1.  (a)  Note  the  physical  properties  of  mercury.     Place  a 
globule  of  it  in  a  small  beaker  and  add  (hood)  just  enough 
nitric  acid  to  dissolve  it.    Dilute  the  product  with  10  cc.  of 
water  and  place  a  copper  penny  in  the  solution.   After  a  few 
minutes  remove  the  coin  and  polish  it  with  a  piece  of  cloth. 
(b)  Account  for  the  result  (recall  the  displacement  series  of 
the  metals). 

2.  For  what  purpose  have  we  used  mercuric  oxide?    Place 
0.5  g.  of  it  in  a  test  tube  and  dissolve  it  in  as  little  nitric  acid 
as  possible,    (c)  What  compound  of  mercury  is  formed?    Then 
add  water  until  the  test  tube  is  one  fourth  full.    Into  a  second 
test  tube  pour  a  similar  volume  of  a  solution  of  mercurous 
nitrate.    Now  add  2  or  3  drops  of  hydrochloric  acid  to  each  test 
tube,    (d)  What  conclusions  do  you  draw  in  reference  to  the 
solubility  of  the  two  chlorides  of  mercury? 

[ISO] 


EXERCISE  101 

A  STUDY  OF  SILVER  AND  ITS  COMPOUNDS 

Apparatus.  2Oo-cc.  beaker;  stirring-rod;  funnel;  blowpipe  and  piece 
of  charcoal;  3  test  tubes. 

Materials.  Silver  dime;  nitric  acid;  hydrochloric  acid;  ammonium 
hydroxide ;  filter  paper ;  hot  water ;  2  or  3  g.  sodium  carbonate ;  10  cc. 
silver  nitrate  solution  (R.  S.) ;  solutions  of  potassium  bromide  and  of 
potassium  iodide  (R.  S.). 

1.  Place  a  silver  dime  in  a  small  beaker  and  add  (hood)  suf- 
ficient nitric  acid  to  dissolve  it.    The  solution  may  be  hastened 
by  applying  a  gentle  heat.   When  the  solution  is  complete, 
dilute  the  product  with  about  25  cc.  of  water,    (a)  Account  for 
the  color  of  the  liquid.    Now  add,  drop  by  drop,  with  con- 
stant stirring,  a  solution  of  hydrochloric  acid.   The  precipitate 
formed  settles  to  the  bottom  of  the  beaker  when  stirred.    Con- 
tinue adding  the  acid  until  a  drop  of  the  acid  no  longer  causes 
a  precipitate  when  it  comes  in  contact  with  the  clear  liquid  in 
the  beaker.    Now  carefully  pour  off  the  clear  liquid  from  the 
precipitate  and  add  ammonium  hydroxide  to  this  liquid  until 
the  solution  becomes  alkaline,    (b)  Account  for  the  change  in 
color  (3,  Exercise  99).   Wash  the  precipitate  remaining  in  the 
beaker  two  or  three  times  by  pouring  hot  water  over  it  and 
decanting.    Finally,  remove  any  remaining  water  by  filtration. 
Mix  the  product  with  an  equal  bulk  of  sodium  carbonate, 
transfer  to  a  small  cavity  in  a  piece  of  charcoal,  and  heat  it  in 
the  blowpipe  flame.    The  silver  salt  is  gradually  reduced  to 
metallic  silver,  which  will  fuse  into  a  globule  if  sufficient  heat 
is  applied,    (c)  How  does  the  product  differ  in  composition 
from  that  of  the  original  coin  ? 

2.  Prepare  small  amounts  of  the  chloride,  the  bromide,  and 
the  iodide  of  silver.    (G?)  Give  the  methods  you  employed.    Ex- 
pose to  the  sunlight  the  test  tubes  containing  the  precipitates 
and  note  any  changes.    For  what  are  these  compounds  used  ? 

F1511 


EXERCISE  102 
THE  CHEMISTRY  OF  PHOTOGRAPHY 

Apparatus.  4  test  tubes;  hard-glass  test  tube;  5oo-cc.  beaker;  glass 
rod ;  piece  of  window  glass. 

Materials.  Solution  of  potassium  bromide  (R.S.);  solution  of  silver 
nitrate  (R.S.);  0.5  g.  silver  nitrate;  0.5  g.  potassium  bromide;  3  g. 
powdered  gelatin;  tube  of  any  commercial  photographic  developer; 
solution  of  sodium  thiosulfate  (R.S.). 

1.  Preparation  of  the  plate  or  film.  Pour  into  a  test  tube  3  cc. 
of  a  solution  of  potassium  bromide  and  add  to  this  2  cc.  of  a 
solution  of  silver  nitrate,  (a)  What  is  the  compound  formed? 
Place  the  tube  and  contents  in  the 
bright  sunlight  for  two  or  three  min- 
utes, (b)  Do  you  observe  any  change 
in  color? 

It  is  not  possible  to  coat  a  glass  plate 
or  film  with  the  pure  silver  bromide. 
The  bromide  is  therefore  suspended  in 
a  finely  divided  state  in  an  emulsion  of 
gelatin,  and  the  glass  coated  with  the 
emulsion.  The  gelatin  keeps  the  bro- 
mide in  a  finely  divided  state  and  also 
makes  it  much  more  sensitive  to  light. 
Proceed  as  follows: 

Half  fill  a  500-cc.  beaker  with  water, 
heat  until  the  water  is  uncomfortably 
hot  to  your  finger,  and  maintain  this  temperature  with  a 
small  flame.  Now  pour  3  cc.  of  water  into  a  test  tube  and 
add  to  it  0.5  g.  of  silver  nitrate.  Shake  the  mixture  until 
the  nitrate  is  dissolved ;  then  place  the  tube  in  the  beaker, 
as  shown  in  Fig.  88.  Next  pour  3  cc.  of  water  in  your 
hard-glass  test  tube  (used  in  preparing  oxygen),  and  place 

[1521 


FIG.  88.  Heating  a  liquid  in 

a  test  tube  by  immersing 

the  tube  in  hot  water 


this  in  the  hot  water  in  the  beaker  until  the  water  in  the  tube 
has  about  the  same  temperature  as  that  in  the  beaker*  Next 
add  to  the  tube  0.5  g.  of  potassium  bromide  and  then,  a  little 
at  a  time,  with  constant  stirring  (use  a  glass  rod),  from  2  to 
3  g.  of  powdered  gelatin.  When  the  gelatin  is  uniformly  dis- 
tributed through  the  hot  water,  add  to  this,  a  little  at  a  time, 
the  hot  solution  of  silver  nitrate  from  the  tube  in  the  beaker, 
stirring  it  thoroughly  with  a  glass  rod.  The  hot  emulsion 
gradually  becomes  thinner.  Now  pour  this  emulsion  (keep  it 
away  from  a  bright  light)  onto  a  glass  plate,  tilting  the  plate 
until  the  emulsion  forms  a  uniform  coating  on  the  plate.  Set 
this  aside  for  five  or  ten  minutes  until  the  emulsion  hardens.  In 
actual  practice  all  these  operations  must  be  carried  out  in  a 
dark  room,  (c)  What  is  the  composition  of  the  emulsion? 
(d)  Note  its  color.  Now  expose  the  plate  to  the  bright  sun- 
light for  a  few  minutes,  (e)  Note  the  change  in  color. 

2.  The  chemistry  of  the  development  of  image.    Pour  into  a 
test  tube  2  cc.  of  a  silver  nitrate  solution  and  add  to  this  3  cc. 
of  potassium  bromide  solution.    Do  not  shake  the  tube,  as  this 
will  cause  the  precipitate  to  form  a  compact  mass.    (/)  Record 
your  results.    Expose  the  tube  and  contents  to  the  bright  light 
for  five  or  ten  seconds ;  then  add  to  it  5  cc.  of  a  solution  of  the 
developer  prepared  as  directed  in  the  instructions  accompany- 
ing the  developer,    (g)  Record  and  explain  the  results.    Save 
the  tube  and  contents  for  experiment  3. 

3.  The  chemistry  of  "fixing"  the  negatives.    Pour  off  the 
liquid  from  the  tube  (experiment  2),  leaving  the  solid  in  the 
tube,  and  add  to  this  solid  10  cc.  of  a  solution  of  sodium  thio- 
sulfate  (hypo).    Shake  the  mixture  vigorously  for  about  one 
minute,    (h)  Does  the  precipitate  dissolve?    Set  the  tube  aside 
for  comparison  later. 

4.  Now  repeat  experiments  2  and  3  in  a  dark  room  (omitting, 
of  course,  the  exposure  to  light,  as  directed  in  experiment  2). 
(i)  Account  for  the  difference  in  the  results  obtained  when  the 
experiment  is  performed  in  the  absence  of  light.    (If  a  dark 

[153] 


room  is  not  available,  fairly  acceptable  results  may  be  obtained 
by  wrapping  a  piece  of  black  paper  around  the  tube  in  which 
the  reactions  are  carried  out  and  performing  the  experiment 
in  a  dark  place  in  the  laboratory.) 

EXERCISE  103 

SOME  PROPERTIES  OF  TIN 

Apparatus.   loo-cc.  beaker;  blowpipe  and  charcoal. 
Materials.    2  g.  tin;  hydrochloric  acid;  solution  of  mercuric  chloride 
(R..S.). 

1.  (a)  Note  the  physical  properties  of  tin.   (jb)  Heat  a  bit  of 
it  on  charcoal  (Fig.  86)  and  note  the  changes. 

2.  Dissolve  about  i  g.  of  the  metal  in  hydrochloric  acid. 
(c)  What  is  formed?    Cool,  dilute  with  a  little  water,  and  add 
i  to  2  drops  of  the  solution  to  3  cc.  of  mercuric  chloride  solution. 
A  white  precipitate  of  mercurous  chloride  forms : 

SnCl2  +  2  HgCl2 >-  SnCl4  +  2  HgCl 

Now  add  a  few  drops  more  of  the  stannous  chloride  solution 
and  heat  the  mixture  gently.  The  mercurous  chloride  is  reduced 
to  metallic  mercury,  which  forms  a  dark-gray  precipitate: 

SnCl2  +  2  HgCl  — >-  SnCl4  +  2  Hg 

EXERCISE  104 

A  STUDY  OF  LEAD  AND  SOME  OF  ITS  COMPOUNDS 

Apparatus.    Blowpipe  and  charcoal ;  5  test  tubes. 

Materials.  2  g.  lead  (obtain  some  scrap  lead  from  a  plumber) ;  lead 
acetate  (R.S.);  0.5  g.  lead  nitrate  dissolved  in  5  cc.  water;  ammo- 
nium sulfide  (R.S.);  sulfuric  acid;  potassium  chromate  (R.S.);  hydro- 
chloric acid  ;  piece  of  mossy  zinc  ;  thread  or  string  about  50  cm.  long ; 
10  cc.  hydrogen  peroxide. 

1.  (a)  Note  the  physical  properties  of  the  metal.  Heat  a 
small  bit  on  charcoal,  (b)  Is  it  easily  melted?  (c)  Account  for 
the  coating  formed  on  the  charcoal. 

[  1541 


2.  To  a  piece  of  mossy  zinc  about  as  large  as  the  end  of  your 
finger  tie  a  piece  of  thread  and  suspend  the  zinc  about  the 
middle  of  a  test  tube  which  is  nearly  filled  with  a  solution  of 
lead  acetate.    Set  the  tube  aside  until  the  end  of  the  laboratory 
period,  noting,  every  few  minutes,  any  changes  taking  place. 
Leave  the  tube  and  contents  until  the  next  laboratory  period 
and  again  note  the  changes,    (d)  Describe  the  results  and  the 
chemistry  involved. 

3.  Pour  i  cc.  of  the  solution  of  lead  nitrate  into  a  test  tube 
and  add  one  drop  of  ammonium  sulfide.   Black  lead  sulfide  is 
formed,    (e)  Write  the  equation  for  the  reaction.  Now  add  to 
the  mixture  7  or  8  cc.  of  hydrogen  peroxide  solution,  shake  the 
mixture,  and  set  it  aside  for  ten  minutes.    (/)  Account  for  the 
change  in  color  (recall  that  hydrogen  peroxide  is  a  good  oxi- 
dizing agent). 

4.  To  the  remainder  of  the  lead  nitrate  solution  add  2  or  3 
drops  of  potassium  chromate  solution.   Yellow  lead  chromate 
(PbCrO4)  is  formed.  This  is  used  as  a  pigment  under  the  name 
of  chrome  yellow. 

Pb(NO3)2  +  K2CrO4 >-  PbCrO4  +  2  KNO3 

EXERCISE  105 

SIMPLE  CELLS  FOR  PRODUCING  ELECTRIC  CURRENTS 

Apparatus.    Beaker  (ioo-cc.);  screw  clamp. 

Materials,  i  strip  each  of  sheet  iron  and  of  zinc  about  i  cm.  x  10  cm. 
50  cc.  water,  to  which  is  added  i  cc.  sulfuric  acid. 

1.  Action  of  metals  in  a  cell  (for  producing  an  electric  cur- 
rent). A  cell  in  its  simplest  form  consists  of  two  strips  of 
different  metals  joined  together  at  one  end  by  a  wire  while  the 
other  ends  are  immersed  in  a  dilute  solution  of  an  acid,  such  as 
sulfuric.  One  of  the  strips  will  slowly  dissolve  in  the  acid  while 
hydrogen  escapes  from  the  surface  of  the  other  metal.  As  a 
result  of  the  action  an  electric  current  is  generated  and  flows 

[  1551 


through  the  wire.    To  determine  which  of  the  metals  is  dis- 
solved, proceed  as  follows : 

Cut  a  strip  of  sheet  iron  about  10  cm.  x  i  cm. ;  also  a  similar 
strip  of  zinc.  Polish  one  end  of  each  strip  so  as  to  remove  all 
foreign  matter  and  then  clamp  the  ends 
firmly  together  by  a  screw  clamp.  Bend  the 
two  strips  in  the  form  of  the  letter  V  and 
immerse  the  disconnected  ends  in  a  solution 
of  sulfuric  acid  (made  by  adding  i  cc.  of 
the  acid  to  50  cc.  of  water),  as  shown  in 
Fig.  89.  (a)  Note  from  which  metal  hydro- 
gen escapes  (if  the  metals  are  impure,  hy- 
drogen may  escape  from  both  metals,  but 
the  amount  escaping  from  one  of  the  metals 
will  always  be  in  excess).  Set  the  beaker 
and  metals  aside  for  one  or  two  hours  (or 
until  the  next  period),  (b)  Note  which  of 
the  metals  has  been  dissolved,  (c)  Note  the  relative  positions 
of  zinc  and  iron  in  the  displacement  series. 

When  any  two  metals  are  connected,  as  in  the  above  experi- 
ment, the  one  nearest  the  top  of  the  displacement  series  is 
always  the  one  which  is  dissolved. 

2.  Galvanized  iron  consists  of  strips  of  iron  coated  with  zinc. 
If  the  zinc  is  worn  away  at  any  point  so  as  to  expose  the 
iron,  and  moisture  is  present,  a  local  current  will  set  up.  (d) 
Which  of  the  two  metals  will  be  dissolved?  (e)  What  is  the 
composition  of  tin  plate?  (/)  If  the  outside  metal  is  worn 
away,  and  moisture  is  present,  which  of  the  two  metals  will 
be  dissolved  ? 


FIG.  89.  A  simple  cell 
for  producing  an  elec- 
tric current 


[156] 


EXERCISE  106 
PAINTS 

Apparatus.  Blowpipe  and  charcoal;  test  tube;  funnel  and  filter 
paper ;  small  beaker. 

Materials.  About  5  g.  of  samples  of  white  paints;  5  cc.  carbon 
tetrachloride;  i  cc.  ammonium  sulfide  (R.  S.);  filter  paper;  splint. 

1.  Introduce  2  or  3  g.  of  a  sample  of  a  white  paint  into  a 
test  tube  and  then  add  carbon  tetrachloride  until  the  tube  is 
about  one  fourth  full.    Shake  the  tube  and  contents  vigorously ; 
then  set  it  aside  for  a  few  minutes  (or  until  you  have  performed 
the  experiments  below),  (a)  Does  the  solid  matter  in  the  paint 
settle  to  the  bottom  of  the  tube?   Shake  the  tube  again,  and 
filter,    (b)  Is  the  filtrate  clear?    (c)  In  what  form  is  the  solid 
matter  present  in  the  paint?  (d)  What  constituent  of  paint 
is  soluble  in  carbon  tetrachloride?   Place  2  or  3  drops  of  the 
filtrate  on  a  piece  of  filter  paper  and  put  it  aside  until  the  car- 
bon tetrachloride  evaporates,    (e)  Does  a  transparent  "spot" 
remain  ?    (/)  If  so,  what  is  the  cause  of  it  ? 

2.  Stir  the  sample  of  paint  and  then  place  2  or  3  drops  of  it 
on  a  cavity  in  the  piece  of  charcoal.    Heat  this  gently  in  the  oxi- 
dizing flame  of  the  blowpipe  until  any  oil  present  burns  away ; 
then  increase  the  heat  for  two  or  three  minutes.  Any  white  lead 
present  is  reduced  to  metallic  lead ;  while  any  zinc  oxide  pres- 
ent will  form  a  coating  which  is  yellow  while  hot,  but  becomes 
white  on  cooling.  In  this  way  test  samples  for  white  lead  and 
zinc  and  (g)  record  your  results. 

3.  Coat  a  wooden  splint  with  white  paint  and  place  it  in  water 
containing  a  few  drops  of  ammonium  sulfide.    (h)  Interpret 
your  results.    (/)  What  would  be  the  result  of  painting  the  in- 
terior of  a  chemical  laboratory  with  white-lead  paint  ? 


[157] 


EXERCISE  107 
A  STUDY  OF  SOME  OF  THE  COMPOUNDS  OF  MANGANESE 

Apparatus.    6  test  tubes. 

Materials,  o.i  to  0.2  g.  potassium  permanganate  (KMn04);  crystal 
of  ferrous  sulfate;  sulfuric  acid;  ammonium  sulfide  (R.  S.);  ammo- 
nium carbonate  (R.  S.);  sodium  hydroxide;  manganese  chloride  solu- 
tion (R.  S.). 

1.  (a)  Examine  the  physical  properties  of  potassium  perman- 
ganate. Dissolve  about  o.i  g.  of  it  in  5  cc.  of  water.  Add  a 
drop  of  the  solution  to  a  solution  of  the  ferrous  sulfate  contain- 
ing 2  or  3  drops  of  sulfuric  acid.  The  ferrous  sulfate  is  changed 
to  ferric  sulfate,  the  oxygen  in  the  reaction  (see  equation  below) 
coming  from  the  permanganate,  which  is  a  good  oxidizing  agent. 


2  FeS04+H2S04  +  0  -  ^Fe2(SO4)3  +  H2O 

2.  In  potassium  permanganate  the  manganese  acts  as  an  a  jid- 
forming  element.  It  also  acts  as  a  base-forming  element  in 
certain  compounds.  Try  the  action  of  ammonium  sulfide,  am- 
monium carbonate,  and  sodium  hydroxide,  respectively,  on  a 
solution  of  manganese  chloride,  (b)  Describe  the  results  and 
write  the  equations  for  the  reactions  involved. 

EXERCISE  108 
A  STUDY  OF  SOME  OF  THE  COMPOUNDS  OF  CHROMIUM 

Apparatus.     6  test  tubes. 

Materials.  Solution  of  potassium  chromate  (R.S.)  ;  lead  acetate 
solution  (R.  S.);  barium  chloride  solution  (R.  S.)  ;  ammonium  sul- 
fide (R.S.)  5  sodium  carbonate  solution  (R.S.)  ;  sodium  hydroxide; 
chromium  sulfate  solution  (R.  S.). 

1.  (a)  Chromates  and  dichromates.  Write  the  formula  for 
potassium  chromate;  for  potassium  dichromate.  (b}  Is  the 
chromium  an  acid-forming  or  a  base-forming  element  in  these 

[1581 


compounds?   Add  2  or  3  drops  of  sulfuric  acid  to  a  little 
potassium  chromate  solution,    (c)  Explain  the  results. 

2.  Try  the  effect  of  a  solution  of  potassium  chromate  on  a 
solution  of  a  compound  of  lead ;  also  on  a  compound  of  barium. 
(d)  Describe  the  results  and  write  the  equations  for  all  the 
reactions  involved. 

3.  Salts  of  chromium.    Try  the  effect  of  the  following  re- 
agents on  a  solution  of  a  salt  of  chromium :  ammonium  sulfide, 
sodium  carbonate,  sodium  hydroxide,    (e)  Describe  the  results 
and  write  the  equations  for  all  the  reactions. 

EXERCISE  109 

THE  DETECTION  OF  SILVER,  LEAD,  AND  MERCURY  WHEN 
PRESENT  IN  THE  SAME  SOLUTION  (OPTIONAL) 

Apparatus.     Two  beakers;  stirring-rod;  funnel;  test  tubes. 

Materials.  Solutions  of  silver  nitrate  (R.S.),  lead  acetate  (R.S.), 
and  mercurous  nitrate  (R.S.) ;  hydrochloric  acid;  filter  paper;  potassium 
chromate  (R. S.);  ammonium  hydroxide;  nitric  acid;  hot  water;  litmus 
paper  (blue). 

1.  The  detection  of  any  one  metal  becomes  more  complicated 
when  other  metals  are  present  in  the  same  solution.  As  a  rule 
it  is  necessary  so  to  treat  the  mixture  as  to  separate  the  metals 
from  each  other.  The  principle  involved  is  illustrated  in  the 
following  experiment. 

Pour  into  a  beaker  5  cc.  of  the  silver  nitrate  solution  and  2  cc. 
each  of  the  solutions  of  lead  acetate  and  mercurous  nitrate. 
Dilute  with  water  to  about  200  cc.  Now  add  hydrochloric  acid, 
drop  by  drop,  with  constant  stirring.  The  precipitate  formed 
on  stirring  settles  to  the  bottom  of  the  beaker.  Continue  add- 
ing the  acid  until  a  drop  of  it  no  longer  causes  a  precipitate 
when  brought  in  contact  with  the  clear  liquid  in  the  beaker. 
(a)  Write  the  equations  for  the  reactions  involved,  (b)  Re- 
cord the  names  and  formulas  of  the  compounds  constituting  the 
precipitate.  Filter  off  the  precipitate  and  fill  the  filter  paper 

f  1591 


with  boiling  water,  collecting  the  filtrate  in  a  beaker.  The  hot 
water  dissolves  the  lead  chloride  so  that  this  compound  is 
present  in  the  filtrate.  Test  for  it  by  adding  to  the  filtrate  a 
few  drops  of  potassium  chromate  solution.  Yellow  lead  chro- 
mate  is  precipitated. 

It  is  important  that  all  the  lead  chloride  should  be  removed 
from  the  mixture  on  the  filter  paper ;  hence  again  fill  the  filter 
paper  with  hot  water.  Wait  until  it  runs  through,  and  again 
repeat  the  process,  discarding  the  filtrates. 

(c)  What  is  the  residue  on  the  filter  paper?   To  this  residue 
add  2  or  3  cc.  of  ammonium  hydroxide  and  collect  the  filtrate 
in  a  test  tube.   This  filtrate  contains  the  silver  chloride  dis- 
solved from  the  residue  by  the  ammonium  hydroxide.    To 
prove  its  presence,  neutralize  the  ammonium  hydroxide  present 
by  adding  nitric  acid  to  the  liquid  until  just  acid  to  litmus 
paper.    The  silver  chloride  is  precipitated. 

(d)  What  effect  did  the  ammonium  hydroxide  have  upon  the 
color  of  the  residue  on  the  filter  paper?    This  change  in  color 
is  due  to  the  action  of  ammonium  hydroxide  on  the  mercurous 
chloride,  and  serves  as  a  test  for  the  presence  of  the  latter. 

(e)  Supposing  that  the  original  solution  contained  only  one 
or  two  of  the  metals  of  the  group,  how  would  the  absence  of 
the  remaining  ones  be  indicated  ? 

EXERCISE  110 

BORAX-BEAD  TESTS  (OPTIONAL) 

Apparatus.    Platinum  wire. 

Materials.  Borax  (R.  S.) ;  small  piece  (size  of  a  pin's  head)  of  a 
compound  of  each  of  the  following  metals :  nickel,  iron,  manganese, 
copper. 

1.  Recall  the  effect  of  adding  a  trace  of  cobalt  nitrate  to  a 
borax  bead  (Exercise  72).  Repeat  the  experiment,  substi- 
tuting for  the  cobalt  nitrate,  salts  of  the  following  metals: 
nickel,  iron,  manganese,  copper,  (a)  Record  your  results. 

1.160] 


APPENDIX 

THE  METRIC  SYSTEM 

With  four  or  five  exceptions  this  system  is  now  used  in  all  civil- 
ized countries.  The  United  States  and  Great  Britain  are  among 
the  few  countries  that  have  not  formally  adopted  it,  but  even  in 
these  countries  the  system  is  universally  used  by  scientists  and  is 
coming  into  use  more  and  more  by  manufacturers. 

In  the  metric  system  each  unit  is  ten  times  as  large  as  the 
next  lower  unit;  hence  the  system  is  often  termed  the  "decimal 
system." 

1.  Length.    The  unit  is  the  meter.    It  is  equal  to  39.37  inches. 

10  millimeters  (mm.)  =  i  centimeter  (cm.) 

10  centimeters  =  i  decimeter  (dm.) 

10  decimeters  =  i  meter  (m.) 

1000  meters  =  i  kilometer  (km.) 

The  only  measures  of  length  ordinarily  used  by  the  chemist  are 
the  millimeter  and  the  centimeter ;  thus,  the  height  of  the  barome- 
ter at  the  sea  level  is  recorded  as  76  cm.  (or  more  commonly  as 
760  mm.)  and  not  7  dm.  and  6  cm. 

2.  Volume.    The  unit  generally  used  is  the  cubic  centimeter. 

1000  cubic  millimeters  =  i  cubic  centimeter  (cc.) 
1000  cubic  centimeters  =  i  cubic  decimeter  =  i  liter 
1000  cubic  decimeters  =  i  cubic  meter 

The  chemist  uses  only  the  cubic  centimeter  and  the  liter  as 
measures  of  volume.  Thus,  the  volume  of  a  test  tube  is  given  as 
(say)  25  cc.;  that  of  a  flask  as  (say)  500  cc.,  or  %  liter. 

[1611 


3.  Weight.  The  unit  is  the  gram.  This  is  approximately  the 
weight  of  i  cc.  of  pure  water  at  its  temperature  of  greatest  den- 
sity (4°).  It  is  equal  to  15.43  grains. 

* 

10  milligrams  (mg.)  =  i  centigram  (eg.) 
10  centigrams  =  i  decigram  (dg.) 
10  decigrams  =  i  gram  (g.) 
1000  grams  =  i  kilogram  (kg.) 

The  gram  and  kilogram  are  the  units  of  weight  most  generally 
used  by  the  chemist.  Thus,  the  weight  of  a  crucible  is  given  as 
(say)  10.532  g.  and  not  10,532  mg.  or  10  g.  5  dg.  3  eg.  2  mg. 

RELATION  BETWEEN  ENGLISH  AND  METRIC  CONSTANTS 

i  pound  (troy)  =  373.24  grams 
i  ounce  (troy)  =  31.10348  grams 
i  pound  (avoirdupois)  =  453.59  grams 
i  ounce  (avoirdupois)  =  28.3495  grams 

i  kilogram  =  2.67923  pounds  (troy) 

i  kilogram  =  2.20462  pounds  (avoirdupois) 

i  liter  =  1.05668  United  States  quarts  ^ 
i  gallon  =  3.78543  liters 
i  cubic  centimeter  =  0.06 10  cubic  inch 

i  cubic  inch  =16.3872  cubic  centimeters 
i  cubic  foot  =  28,320  cubic  centimeters 
i  centimeter  =  0.3937  inch 
i  meter  =  39.37  inched 


Also  note  that 


i  centimeter  =  nearly  f  inch 
.      i  meter  =  nearly  i.i  yards 
i  kilogram  =  nearly  2\  pounds  avoirdupois 


[  162 


TABLE  OF  SOLUBILITIES  OF  SOME  OF  THE  COMPOUNDS 
OF  THE  METALS 


ACETATE 

BROMIDE 

CARBONATE 

CHLORATE 

CHLORIDE 

CHROMATE 

HYDROXIDE 

IODIDE 

NITRATE 

§ 
1 

PHOSPHATE 

SILICATE  (ORTHO) 

SULFATE 

SULFIDE 

Aluminium  .... 

w 

w 

W 

W 

A 

W 

W 

A 

A 

A 

W 

A 

Ammonium      .     .     . 

w 

w 

W 

W 

W 

W 

W 

w 

w 

W 

w 

W 

Barium   

w 

w 

A 

W 

W 

A 

W 

w 

w 

A 

A 

A 

I 

W 

Calcium  

w 

w 

A 

W 

W 

X 

W 

w 

w 

X 

A 

A 

X 

X 

Cobalt     

w 

w 

A 

w 

w 

A 

A 

w 

w 

A 

A 

A 

w 

A 

Copper 

w 

w 

A 

w 

w 

W 

A 

w 

w 

A 

A 

A 

w 

A 

Ferric 

w 

w 

w 

w 

W 

A 

w 

w 

A 

A 

A 

w 

A 

Ferrous  

w 

w 

A 

w 

w 

A 

w 

w 

A 

A 

A 

w 

A 

Lead  

w 

X 

A 

w 

X 

A 

A 

X 

w 

A 

A 

A 

I 

A 

Magnesium 

w 

w 

A 

w 

w 

W 

A 

w 

w 

A 

A 

A 

w 

A 

Manganese 

w 

w 

A 

w 

w 

W 

A 

w 

w 

A 

A 

A 

w 

A 

Mercuric     .... 

w 

w 

A 

w 

w 

X 

A 

w 

A 

A 

X 

I 

Mercurous  .... 

w 

A 

w 

A 

A 

A 

w 

A 

A 

X 

Nickel    

w 

w 

A 

w 

W 

A 

A 

W 

w 

A 

A 

A 

w 

A 

Potassium    .... 

w 

w 

W 

w 

W 

W 

W 

W 

w 

W 

W 

W 

w 

W 

Silver      

w 

I 

A 

w 

I 

A 

I 

w 

A 

A 

X 

A 

Sodium  

w 

w 

W 

w 

w 

W 

W 

W 

w 

W 

W 

W 

w 

W 

Stannic  

w 

w 

w 

A 

W 

A 

A 

A 

Stannous      .... 

w 

w 

w 

A 

A 

w 

A 

A 

w 

A 

Zinc  .      ..... 

w 

w 

A 

w 

w 

W 

A 

w 

w 

A 

A 

A 

w 

A 

W,  soluble  in  water. 

A,  insoluble  in  water;  soluble  in  either  HC1  or  HNO3  or  in  both. 
I,  insoluble  in  water  and  in  acids. 

X,  slightly  soluble  in  water  and  slightly  or  readily  soluble  in  acids. 

[  1631 


TREATMENT  IN  CASE  OF  ACCIDENT 

Every  laboratory  should  be  supplied  with  the  materials  neces- 
sary for  the  treatment  of  cuts  and  burns.  Such  wounds  when  im- 
mediately and  intelligently  treated  give  little  or  no  trouble,  but 
if  left  to  take  care  of  themselves  infection  may  occur  and  serious 
results  follow.  In  case  of  severe  wounds  the  student,  after  the  pre- 
liminary treatment  in  the  laboratory,  should  be  sent  to  a  physician. 

The  following  materials  required  for  the  treatment  of  wounds 
should  be  kept  in  the  laboratory  in  a  tight  cabinet  or  cupboard : 

1.  Plain,  sterile  gauze  bandage  :  6  rolls,  i  in.  x  10  yd.;  6  rolls,  2  in.  x 
10  yd. 

2.  Adhesive  plaster  for  holding  gauze  in  place  :  2  rolls,  0.5  in.  x  10  yd. 

3.  1000  cc.  of  benzine  (low-boiling  gasoline). 

4.  2000  cc.  of  Seiler's  solution  (this  may  be  made  from  tablets  pur- 
chased at  any  drug  store). 

5.  Boric  acid  solution  prepared  by  dissolving  20  g.  of  boric  acid  in 
1000  cc.  of  water. 

6.  Iodine  solution — ordinary  tincture  of  iodine   (this  may  be  pur- 
chased at  any  drug  store). 

7.  Agnew's  solution  prepared  after  the  following  prescription: 

Tannic  acid 2  grams 

Borax 2  grams 

Glycerin n  milliliters 

Camphorated  water      .      .     .  q.  s.  to  make  90  milliliters 

Different  wounds  should  be  treated  as  follows : 

1.  Cuts  from  glass.      Remove  any  dirt  or  grease  from  the  wound 
and  surrounding  skin  by  washing  with  gauze  saturated  with  benzine. 
A  rather  free  bleeding  at  first  will  help  to  prevent  infection.    Finally, 
wash  the  wound  with  a  piece  of  gauze  saturated  with  the  iodine  solu- 
tion; then  bandage  so  as  to  prevent  contamination. 

2.  Burns.     Immediately  apply  gauze  saturated  with  Seiler's  solution 
and  then  bandage. 

3.  Burns  from  acids.     Treat  the  same  as  ordinary  burns,  as  directed 
under  2. 

4.  Burns  from  alkalies.   Immediately  wash  with  a  large  volume  of 
water  and  then  with  boric  acid  solution  on  gauze ;  then  bandage. 

[164] 


5.  Acid  in  eyes.  Immediately  wash  the  eyes  with  a  large  volume  of 
water,  then  with  a  solution  of  boric  acid.  Finally,  drop  into  the  eyes 
3  or  4  drops  of  Agnew's  solution  and  then  apply  to  the  eyes  a  small  gauze 
pad  which  is  kept  saturated  with  ice  water. 

INFORMATION  REGARDING  APPARATUS  AND  CHEMICALS 

The  lists  following  include  the  apparatus  and  chemicals  required 
for  the  experiments  in  this  notebook.  It  is  always  best  to  furnish 
each  student  with  as  complete  an  outfit  as  possible  and  to  hold  him 
responsible  for  it.  Certain  pieces  may,  however,  be  used  in  com- 
mon by  a  number  of  students,  and  these  have  been  placed  in  a 
separate  list.  It  is  always  cheapest  to  purchase  the  apparatus  and 
chemicals  in  as  large  quantities  as  possible.  The  amounts  of  most 
of  the  chemicals  needed  for  a  class  of  ten  are  so  small  that  their 
cost  will  be  proportionately  much  greater  than  when  larger  quan- 
tities are  ordered.  It  is  always  best  to  order  the  definite  amounts 
of  chemicals  listed  in  the  catalogues,  such  as  100  g.  or  i  lb.; 
otherwise  the  cost  of  weighing  out  odd  quantities  and  preparing 
these  for  shipment  may  amount  to  more  than  the  cost  of  the 
chemicals.  The  supplies  may  be  obtained  from  any  of  the  large 
dealers.  Catalogues  will  be  sent  on  application  and  should  be  in 
every  school.  The  following  are  the  addresses  of  some  of  the 
largest  firms:  Burrell  Technical  Supply  Co.,  Pittsburgh,  Pa.; 
Central  Scientific  Company,  Chicago,  111. ;  The  Chemical  Rubber 
Co.,  Cleveland,  Ohio;  Chicago  Apparatus  Co.,  Chicago.  111.; 
Kauffman-Lattimer  Co.,  Columbus,  Ohio ;  L.  E.  Knott  Apparatus 
Co.,  Boston,  Mass. ;  Eimer  and  Amend,  New  York  City ;  Standard 
Scientific  Co.,  New  York  City;  Scientific  Materials  Co.,  Pitts- 
burgh, Pa. ;  E.  H.  Sargent  &  Co.,  Chicago,  111. ;  Arthur  H.  Thomas 
Company,  Philadelphia,  Pa.;  W.  M.  Welch  Mfg.  Co.,  Chicago, 
111. ;  The  Will  Corporation,  Rochester,  New  York. 

A  list  of  the  supplies  needed  should  be  sent  to  a  number  of 
firms  for  quotations  on  prices.  In  ordering  any  piece  of  apparatus 
a  certain  form  in  some  catalogue  should  be  designated;  other- 
wise it  will  be  impossible  to  compare  the  prices.  In  general  it  is 
best  to  purchase  as  simple  a  form  of  apparatus  as  possible;  for 
example,  20  cents  will  buy  a  Bunsen  burner  which  for  ordinary 

[165] 


purposes  is  preferable  to  those  costing  $i.  A  person  experienced 
in  the  purchase  of  supplies  will  always  find  it  possible  to  reduce 
materially  the  cost  of  the  order. 

It  is  generally  cheaper  to  purchase  from  local  stores  ordinary 
pieces  of  apparatus  and  materials  commonly  sold  by  such  stores, 
rather  than  to  order  them  through  supply  houses.  Thus,  an  iron 
spoon  can  be  purchased  at  any  local  five-and-ten-cent  store  and 
will  cost  less  than  when  purchased  through  a  supply  house.  Simi- 
larly, common  salt  can  be  obtained  at  a  lower  cost  at  the  corner 
grocery  than  through  a  supply  house.  A  few  substances  are  listed 
that  are  not  ordinarily  furnished  by  supply  houses.  Thus,  such 
houses  do  not  ordinarily  list  "natural  rubber,"  but  this  can  be 
purchased  at  a  trifling  cost  from  any  manufacturer  of  rubber  tires 
or  other  rubber  commodities.  Denatured  alcohol  may  be  used  where 
alcohol  is  specified,  but  it  is  better  to  use  the  pure  alcohol.  This 
can  be  secured  tax  free  from  supply  houses  by  presenting  a  per- 
mit secured  from  the  Federal  prohibition  officer  of  your  state.  Such 
alcohol  must  be  kept  by  the  instructor  and  given  out  to  students 
only  on  signed  blanks  giving  the  use  for  which  the  alcohol  is  desired. 

Materials  purchased  of  supply  houses  should  be  ordered  at 
least  three  or  four  months  in  advance,  for  the  rate  of  delivery  is 
often  very  slow.  It  is  best  for  the  teacher,  during  the  spring 
months,  to  make  up  and  place  the  entire  order  for  the  following 
academic  year.  The  teacher  should  be  careful  in  giving  exact 
specifications  and  should  insist  on  getting  what  he  orders.  All 
materials  not  in  accord  with  the  specifications  should  be  promptly 
returned;  otherwise  his  laboratory  may  become  the  dumping 
ground  for  poor  or  off-sized  materials.  It  makes  a  wide  difference 
in  the  student's  work  whether  or  not  he  is  compelled  to  use  (say) 
glass  tubing  that  is  off  size  and  corks  that  do  not  fit.  The  right  size 
costs  no  more.  Be  sure  to  specify  it  and  then  insist  on  getting  it. 

APPARATUS  REQUIRED  FOR  EACH  STUDENT  (TO  BE  KEPT  IN 

STUDENT'S  LOCKER) 

Beakers,  nest  of  seven,  from  loo-cc.  to  yoo-cc. 

Blowpipe,  mouth. 

Bottles,  wide-mouthed:  one,  6o-cc.;  five,  250-00. 

[  1661 


Burner,  wing-top,  for  bending  glass  tubing  (Fig.  12). 

Calcium  chloride  drying-tube,  straight,  15  cm.  in  length  (Fig.  37,  B). 

Charcoal,  i  piece  about  8  cm.  x  3  cm.  x  2  cm. 

Copper  gauze,  10  cm.  x  12  cm.,  from  60  to  70  mesh. 

Deflagra  ting-spoon. 

Dish,  lead,  diameter  about  6  cm.,  depth  3  cm. 

Evaporating-dish,  diameter  7  cm. 

Files  :  round,  about  15  cm.  in  length ;  triangular,  about  15  cm.  in  length. 

Filters,  25,  diameter  about  n  cm. 

Florence  flasks:  two,  25o-cc.;  one,  soo-cc. 

Forceps  (steel),  10  cm.  long. 

Funnel,  diameter  about  6.5  cm. 

Funnel  tube,  external  diameter  of  tube  6  mm. 

Glass  tubing :  300  g.,  soft,  external  diameter  6  mm.,  walls  i  mm.  thick; 
i  piece,  hard-glass,  30  cm.  in  length,  internal  diameter  1.5  cm., 
walls  1.8  mm.  thick;  i  piece,  hard  glass,  25  cm.  in  length  and 
6  mm.  internal  diameter. 

Glass  rod,  diameter  2  mm.,  i  piece. 30  cm.  in  length. 

Medicine  dropper. 

Mortar  (diameter  about  8  cm.)  and  pestle  (both  of  porcelain). 

Pipestem  triangle  for  holding  porcelain  crucible  (Fig.  30). 

Platinum  wire,  small  (No.  28),  5  cm.  long,  for  flame  tests. 

Porcelain  crucible  and  lid,  diameter  about  3.5  cm. 

Rubber  tubing:  internal  diameter  5  mm.,  i  piece  50  cm.  in  length; 
soft,  pure  gum,  internal  diameter  5  mm.,  i  piece  30  cm.  in  length 
for  connections  etc. 

Screw  clamp  for  rubber  tubing. 

Splints  (ordinary  cigar  lighters),  125. 

Sponge. 

Spoon,  ordinary  iron,  bowl  5  or  6  cm.  long. 

Stoppers :  rubber,  one-hole,  to  fit  the  hard-glass  test  tube ;  rubber,  two- 
hole,  to  fit  wide-mouthed  25o-cc.  bottle;  rubber,  two-hole,  to  fit 
250-cc.  flask;  rubber,  two-hole,  to  fit  looo-cc.  narrow-mouthed 
bottle. 

Stoppers,  rubber,  2  one-hole,  to  fit  the  hard-glass  tubing  specified  above. 

Test  tube,  graduated,  3o-cc.,  about  20  cm.  long,  with  o.5-cc.  graduations. 

Test  tube,  hard-glass,  15  cm.  in  length,  diameter  about  1.8  cm. 

Test  tubes,  12,  length  15  cm.,  diameter  about  1.5  cm. 

Test-tube  brush. 

Test-tube  rack. 

Towel. 

[167] 


Tripod. 

Watch  glass,  diameter  about  8  cm. 
Window  glass,  4  pieces  10  cm.  square. 
Wire  gauze,  2  pieces  12  cm.  square. 

APPARATUS  TO  BE  LEFT  ON  EACH  DESK 

Bunsen  burner,  with  75  cm.  of  rubber  tubing  to  fit. 

Clamp,  iron,  large,  for  holding  flasks  and  condensers. 

Iron  tripod  (Fig.  58). 

Pneumatic  trough  (Fig.  28).  The  trough  should  be  about  12  to  15  cm. 
deep,  and  large  enough  to  hold  4  or  5  wide-mouthed  bottles 
(25o-cc.).  It  may  be  round  or  rectangular.  A  pan  made  of  granite 
ware  or  an  earthen  crock  serves  well,  or  any  tinsmith  can  readily 
make  suitable  troughs  of  galvanized  iron. 

Ring  stand  and  3  rings. 

Sand  bath,  iron,  about  12  cm.  in  diameter. 

REAGENTS  ON  EACH  DESK 

250-cc.  bottles  filled  with  the  reagents  named  below.   The  bottle  con- 
taining the  sodium  hydroxide  should  have  ordinary  corks,  the  others 
should  be  glass-stoppered. 
Ammonium  hydroxide  (density  0.90). 
Hydrochloric  acid  (density  1.2). 
Nitric  acid  (density  1.4). 

Sodium  hydroxide  solution  (10  g.  in  100  cc.  of  water). 
Sulfuric  acid  (density  1.84). 

GENERAL  APPARATUS  FOR  TEN  STUDENTS 

Apparatus  marked  with  a  star  (*)  are  required  for  optional  experiments. 

*2  sets  apparatus  for  testing  conductivity  of  solutions  (Fig.  62) ;  this 
may  be  made  or  may  be  purchased  of  supply  houses. 

*i  Babcock  milk-test  apparatus  complete,  either  two-  or  four-bottle 
(Fig.  76). 

1  balance,  weighing  from  0.5  g.  to  500  g.,  with  accompanying  weights. 

2  balances,  sensitive  to  i  eg.  and  made  to  carry  a  load  of  100  g.  with 

accompanying  weights  (each  set  from  i  eg.  to  50  g.  in  covered 

wooden  box), 
i  barometer. 

5  one-liter,  narrow-necked  bottles, 
i  bottle  or  flask,  2ooo-cc.  (Fig.  43,  B). 
4  burettes,  5o-cc.,  graduated  in  o.i  cc.  (Fig.  61). 

[168] 


5  pieces  cobalt  glass  10  cm.  square  (for  flame  tests). 

*4  condensers  (Liebig),  with  rubber  tubing. 

2  sets  cork  borers  (6  in  a  set),  and  sharpener. 

2  gross  corks,  best  grade,  sizes  7,  8,  9,  10,  and  12. 

i  cylinder,  graduated,  loo-cc. 

i  cylinder,  graduated,  5oo-cc. 

i  distilling  apparatus  for  preparing  distilled  water. 

i  magnifying  glass,  small. 

i  microscope,  eyepiece  i  inch,  objectives  f  and  £. 

5  graduated  pipettes  (5  cc.,  graduated  in  divisions  of  o.i  cc.  or  0.2  cc.). 

*i  spectroscope  (Fig.  84). 

5  thermometers,  graduated  from  — 10°  to  150°  C. 

5  retorts,  glass-stoppered,  iso-cc.  (Fig.  65). 

5  porcelain  boats  about  8  cm.  in  length  and  6  or  7  mm.  wide  (these 
must  be  of  such  a  size  as  to  slip  easily  into  the  hard-glass  tub- 
ing, specified  under  the  heading  "Apparatus  required  for  Each 
Student"). 

CHEMICALS  ON  REAGENT  SHELF  (FOR  USE  OF  ALL  STUDENTS) 

The  bottles  containing  solutions  should  be  glass-stoppered. 
Gummed  letters  of  the  alphabet,  of  different  sizes,  may  be  ob- 
tained at  a  little  cost  from  the  Tablet  and  Ticket  Co.,  Chicago, 
111.,  and  these  may  be  used  in  making  the  labels  for  the  bottles ;  or 
the  complete  labels  for  all  the  common  chemicals  may  be  pur- 
chased from  supply  houses.  If  the  class  is  small,  bottles  holding 
250  cc.  will  ordinarily  serve;  if  the  class  is  large,  then  it  is  better 
to  use  bottles  holding  at  least  500  cc.  A  few  of  the  reagents,  such 
as  limewater,  are  used  so  extensively  that  it  is  better  to  use  a 
looo-cc.  bottle.  Distilled  water  must  be  used  in  making  all  solutions. 
A  10  per  cent  solution  signifies  10  g.  dissolved  in  100  cc.  of  water. 

Acetic  acid  (36  per  cent). 

Aluminium  sulfate  (10  per  cent  solution). 

Ammonium  carbonate  (25  g.  of  the  solid  dissolved  in  70  cc.  of  water 

and  10  cc.  of  ammonium  hydroxide  (density  0.90),  and  the  solution 

diluted  to  100  cc.  with  water). 
Ammonium  chloride  (10  per  cent  solution). 

Ammonium  molybdate  solution   (obtained  directly  from  dealers). 
Ammonium  sulfide  solution  (obtained  directly  from  dealers). 
Barium  chloride  (10  per  cent  solution). 

[  1691 


Borax  (solid). 

Calcium  chloride  (10  per  cent  solution). 

Carbon  tetrachloride. 

Chloroform. 

Chlorine  water  (water  saturated  with  chlorine). 

Chromium  sulfate  (10  per  cent  solution). 

Cobalt  nitrate  (5  per  cent  solution). 

Copper  sulfate  (10  per  cent  solution). 

Disodium  phosphate  (10  per  cent  solution). 

Fehling's  solution :  solution  A,  prepared  by  dissolving  17.5  g.  of  copper 
sulfate  crystals  in  250  cc.  of  water;  solution  B,  prepared  by  dis- 
solving 87.5  g.  of  sodium  potassium  tartrate  in  250  cc.  of  10  per  cent 
sodium  hydroxide  solution. 

Ferric  chloride  (10  per  cent  solution). 

Ferric  sulfate  (10  per  cent  solution). 

Hydrogen  sulfide  solution.  Prepared  by  passing  hydrogen  sulfide  slowly 
through  water  until  the  water  is  saturated  with  the  gas.  Since  the 
solution  loses  its  strength  after  a  few  days,  it  is  necessary  to  pre- 
pare it  as  needed. 

Iodine  solution.  Prepared  by  dissolving  2  g.  of  iodine  and  log.  of 
potassium  iodide  in  100  cc.  of  water. 

Lead  acetate  (10  per  cent  solution).  If  the  solution  is  not  clear,  add  a 
few  drops  of  acetic  acid  and  shake  the  mixture ;  repeat  until  the 
solution  becomes  clear. 

Limewater  (saturated  solution  of  calcium  hydroxide).  Shake  3  or  4  g.  of 
the  solid  hydroxide  with  i  liter  of  water ;  then  set  the  mixture  aside 
until  the  excess  of  solid  settles.  Use  the  clear  supernatant  liquid. 

Litmus  solution.  Dissolve  sufficient  litmus  (litmus  cubes)  in  the  water 
to  impart  to  it  a  deep  color. 

Magnesium  sulfate  (10  per  cent  solution). 

Manganese  chloride  (10  per  cent  solution). 

Mercuric  chloride  (5  per  cent  solution). 

Mercurous  nitrate  (10  per  cent  solution). 

Phenolphthalein  (i  g.  dissolved  in  200  cc.  of  alcohol). 

Potassium  bromide  (10  per  cent  solution). 

Potassium  chromate  (10  per  cent  solution). 

Potassium  ferrocyanide  (10  per  cent  solution). 

Potassium  hydroxide  (10  per  cent  solution). 

Potassium  iodide  (10  per  cent  solution). 

Potassium  sulfocyanate  (10  per  cent  solution). 

Silver  nitrate  (4  per  cent  solution). 

[1701 


Starch  solution.  Prepared  by  rubbing  to  a  thin  paste  4  or  5  g.  of  starch 
with  cold  water  and  then  adding,  3  or  4  drops  at  a  time  and  with 
stirring,  to  i  liter  of  boiling  water.  Add  also  about  10  g.  of  zinc 
chloride  (this  acts  as  a  preservative).  Mix  thoroughly,  set  the  mix- 
ture aside,  and  use  the  clear  supernatant  liquid. 

Sodium  carbonate  (10  per  cent  solution). 

Sodium  chloride  (solid). 

Sodium  thiosulfate  (10  per  cent  solution). 

Zinc  acetate  (10  per  cent  solution). 

CHEMICALS  REQUIRED  FOR  A  CLASS  OF  TEN 

The  terms  in  parentheses  after  the  names  of  the  chemicals  refer 
to  the  grade  of  materials  to  be  purchased.  The  abbreviation  c.p. 
signifies  "  chemically  pure."  The  weights  required  in  each  case 
are  given  in  the  metric  system,  but  there  is  also  added  the  approxi- 
mate English  equivalent.  Those  chemicals  marked  with  a  star  (*) 
are  for  optional  experiments. 

This  list  does  not  include  substances  always  easily  obtained  at 
home  stores,  such  as  baking  powders,  butter,  clay,  cloth,  fats  of 
various  kinds,  flour,  gasoline,  gelatin,  ice,  iron  and  copper  wire, 
kerosene,  lead,  milk,  molasses,  oleomargarine,  salt,  soap,  starch, 
sugar,  tacks,  vinegar,  yeast,  etc.  APPROXIMATE  AMOUNTS 

Acid,  acetic  (36  per  cent)  (c.p.) 500  g.  (i  Ib.) 

Acid,  acetic  (glacial) 200  g.  (^  Ib.) 

*Acid,  formic  (50  per  cent) 2Op  g.  (£  Ib.) 

Acid,  hydrochloric  (density  1.2)  (c.p.) 2  kg.  (5  Ib.) 

Acid,  hydrochloric  (commercial)  for  use  in  preparing 

carbon  dioxide 2  kg.  (5  Ib.) 

Acid,  nitric  (density  1.4)  (c.p.) 3  kg.  (7  Ib.) 

*Acid,  oxalic  (pure) 200  g.  Q  Ib.) 

Acid,  sulfuric  (density  1.84)  (c.p.) 4  kg.  (9  Ib.) 

*Acid,  sulfuric  (commercial)  (density  1.83)  for  Bab- 
cock  milk  test 500  g.  (i  Ib.) 

Acid,  tannic    (commercial) 100  g.  (£  Ib.) 

Alcohol  (95  per  cent) i  liter  (i  qt.) 

Alum  (ammonium)  pure 500  g.  (i  Ib.) 

Aluminium  (turnings  or  filings)         30  g.  (i  oz.) 

Aluminium  sulfate  (pure,  crystals) 500  g.  (i  Ib.) 

Ammonium  carbonate  (pure) 100  g.  (i  Ib.) 

[171] 


APPROXIMATE  AMOUNTS 

Ammonium  chloride  (pure) 200  g.  (£  Ib.) 

Ammonium  hydroxide  (density  0.90)  (c.p.)     ...  2  kg.  (5  Ib.) 

Ammonium  molybdate  solution 500  g.  (i  Ib.) 

Ammonium  nitrate  (pure) 100  g.  (^  Ib.) 

Ammonium  sulfate  (commercial) 500  g.  (i  Ib.) 

Ammonium  sulfide  solution 500  g.  (i  Ib.) 

Aniline 100  g.  (3  oz.) 

Antimony 30  g.  (i  oz.) 

Arsenic 30  g.  (i  oz.) 

Arsenic  trioxide  (arsenious  oxide)   (commercial)   .      .  30  g.  (i  oz.) 

Barium  chloride  (c.p.) 100  g.  ($  Ib.) 

Barium  nitrate  (crystals) 100  g.  (£  Ib.) 

Barium  sulfate 500  g.  (i  Ib.) 

Benzene 200  g.  (i  Ib.) 

Bismuth .  30  g.  (i  oz.) 

Bleaching-powder 500  g.  (i  Ib.) 

Boneblack    .     .     .     ....     .     .     .     .     .     .  200  g.  (£  lb.) 

Borax  (commercial) 500  g.  (i  lb.) 

Boric  acid  (pure)  .     .     . 100  g.  (£  lb.) 

Cadmium  chloride  (c.p.) 30  g.  (i  03.) 

Calcium  carbide 500  g.  (i  lb.) 

Calcium  carbonate  (precipitated)      .      .      .      .'    .     .  500  g.  (i  lb.) 
Calcium  chloride  (fused  or  granular)  for  filling  drying 

tubes i  kg.  (2  lb.) 

Calcium  fluoride  (fluorspar) 200  g.  Q  lb.) 

Calcium  hydroxide  (hydra ted  lime) 500  g.  (i  lb.) 

Calcium  sulfate  (plaster  of  Paris) i  kg.  (2  lb.) 

Carbon  disulfide  (commercial) 100  g.  (4  oz.) 

Carbon  tetrachloride  (commercial) 2  kg.  (5  lb.) 

Charcoal  (small  pieces) 500  g.  (i  lb.) 

Chloroform   (commercial) 200  g.  (£  lb.) 

Chromium    sulfate 30  g.  (i  oz.) 

Cobalt  nitrate  (pure) 30  g.  (i  oz.) 

Cobalt  chloride 30  g.  (i  oz.) 

Copper  (turnings  or  scrap) 200  g.  Q  lb.) 

Copper  foil  (thin) 100  g.  (£  lb.) 

Copper  nitrate  (pure) 30  g.  (i  oz.) 

Copper  oxide  (black,  finely  powdered) 100  g.  (^  lb.) 

Copper  sulfate  crystals  (c.p.) 200  g.  (%  lb.) 

Cottonseed  oil 200  g.  (i  lb.) 

[172] 


APPROXIMATE  AMOUNTS 

Cyanamide   (commercial) i  kg.  (2  Ib.) 

Dyes :  Gallein,  fuchsine,  methyl  violet,  malachite  green, 

Congo  red 10  g.  of  each 

alizarin  paste  (20  per  cent)     .           100  g.  (£  Ib.) 

Formalin 100  g.  (£  Ib.) 

Gypsum    (crystals) 200  g.  (i  Ib.) 

Hydrogen  peroxide 200  g.  (£  Ib.) 

Iodine 30  g.  (i  oz.) 

Iron  chloride   (ferric)    (c.p.) 60  g.  (2  oz.) 

Iron  powder  (iron  reduced  by  alcohol) 100  g.  (^  Ib.) 

Iron  sulfate  (ferrous) 200  g.  (J  Ib.) 

Iron  sulfide  (ordinary  lumps  for  preparing  hydrogen 

sulfide) i  kg.  (2  Ib.) 

Iron  wire  (picture-frame  wire),  No.  o 25m.  (25yd.) 

Junket  tablets  (obtained  from  any  grocer  or  druggist)  10  tablets 

Lead  acetate  (sugar  of  lead)  (powdered)    ....  200  g.  Q  Ib.) 

Lead  monoxide  (commercial) 100  g.  (J  Ib.) 

Lead  nitrate  (pure) 60  g.  (2  oz.) 

Litmus  cubes 60  g.  (2  oz.) 

Litmus  paper  (100  strips  red,  100  strips  blue)  ...  2  tubes  of  each 

Magnesium  carbonate  (powdered) 100  g.  (£  Ib.) 

Magnesium  sulfate  (Epsom  salts) 500  g.  (i  Ib.) 

Magnesium  wire  or  ribbon 30  g.  (i  oz.) 

Manganese  chloride 60  g.  (2  oz.) 

Manganese  dioxide  (pure,  powdered) i  kg.  (2  Ib.) 

Marble  (pieces  size  of  a  walnut) 2  kg.  (5  Ib.) 

Mercuric  chloride  (corrosive  sublimate)     .      .      .      .  30  g.  (i  oz.) 

Mercuric  nitrate  (c.p.) 30  g.  (i  oz.) 

Mercuric  oxide 30  g.  (i  oz.) 

Mercurous  nitrate  (c.p.) 100  g.  (i  Ib.) 

Mercury 30  g.  (i  oz.) 

Nickel  nitrate 30  g.  (i  oz.) 

Paraffin 500  g.  (i  Ib.) 

Phenolphthalein 30  g.  (i  oz.) 

Phosphorus  (white) 20  g.  (i  oz.) 

Photographic  developer i  tube 

Portland  cement i  kg.  (2  Ib.) 

Potassium  bitartrate  (cream  of  tartar)      ....  100  g.  (^  Ib.) 

Potassium  bromide  (granular,  pure)      .      .      .      .     .  60  g.  (2  oz.) 

Potassium  carbonate  (c.p.) 60  g.  (2  oz.) 

[173] 


APPROXIMATE  AMOUNTS 

Potassium  chlorate  (small  crystals) 500  g.  (i  Ib.) 

Potassium  chloride   (c.p.) 100  g.  (i  Ib.) 

Potassium  chromate  or  sodium  chromate  (pure)      .     .  100  g.  (i  Ib.) 

Potassium  chromium  sulfate  (chrome  alum)     .     .     .  100  g.  (£  Ib.) 

Potassium  dichromate   (pure) 100  g.  Q  Ib.) 

^Potassium  ferrocyanide  (c.p.) 100  g.  ($  Ib.) 

Potassium  hydroxide  (sticks) 100  g.  (£  Ib.) 

Potassium  iodide  (pure) 60  g.  (2  oz.) 

Potassium  nitrate  (pure) 200  g.  (£  Ib.) 

Potassium  permanganate  (pure) 30  g.  (i  oz.) 

Potassium  sulfate  (pure,  anhydrous) 100  g.  (£  Ib.) 

*Potassium  sulfocyanate  (c.p.)  .     . 30  g.  (i  oz.) 

Rubber  (natural)   .     .     .     .   .  .  Jf .     .-..,.  500  g.  (i  Ib.) 

Silver  nitrate .     .     .     ,;   .  30  g.  (i  oz.) 

Soda  lime  (granular)  .     .     .     -.- -.-. »     .     ,     •<..«-;.,  500  g.  (i  Ib.) 

Sodium ..,.*,  30  g.  (i  oz.) 

*Sodium  acetate  (fused) 500  g.  (i  Ib.) 

*Sodium  benzoate  (pure)      .     .     .     .     .     .     .     .  30  g.  (i  oz.) 

Sodium  bicarbonate  (baking-soda) ;.';,  500  g.  (i  Ib.) 

Sodium  bromide  (crystals)  (pure)    ......  200  g.  (£  )b.) 

Sodium  carbonate  (pure,  anhydrous) 200  g.  (I  Ib.) 

Sodium    hydrogen    phosphate    (disodium    phosphate) 

(c.p.) '.  100  g.  (J  Ib.) 

Sodium  hydroxide  (sticks) .  i  kg.  (2  Ib.) 

Sodium  hydroxide  (normal  solution) i  liter 

Sodium  iodide  (crystals)  (pure) 100  g.  (£  Ib.) 

Sodium  nitrate  (pure) 200  g.  (^  Ib.) 

Sodium  potassium  tartrate  (Rochelle  salts)  (powdered)  500  g.  (i  Ib.) 

Sodium  silicate  solution  (water  glass)     .     .     ....  500  g.  (i  Ib.) 

Sodium  sulfate  (crystals)  (Glauber's  salt)  .     ....  200  g.  (^  Ib.) 

Sodium   thiosulfate .  500  g.  (i  Ib.) 

Sulfur  (flowers) 500  g.  (i  Ib.) 

Sulfur  (roll) 500  g.  (i  Ib.) 

Tartar  emetic  (potassium  antimonyl  tartrate)     .     »  60  g.  (2  oz.) 

Tin   (mossy) ,  60  g.  (2  oz.) 

Zinc  (mossy,  arsenic-free) ;  i  kg.  (2  Ib.) 

Zinc   (sheet) .  200  g.  Q  Ib.) 

Zinc  chloride 100  g.  (£  Ib.) 

Zinc  sulfate  (crystals) ...  100  g.  (i  Ib.) 

*Zinc  sulfide '  .     .     ,  ~  .     .     .  30  g.  (i  oz.) 

[174] 


TABLE  OF  CONSTANTS 


LIST  OF  THE  COMMON  ELEMENTS,  THEIR  SYMBOLS, 
THEIR  ATOMIC  WEIGHTS 

AND 

0  = 

-  16 

Aluminium    . 

.     Al 

27.1 

Iodine  .      .     . 

.     I 

126.92 

Antimony 

.'    Sb 

I2O.2 

Iron      .      .      . 

.     Fe 

55.84 

Arsenic    . 

.      As 

74.96 

Lead     .      .      . 

.     Pb 

207.2 

Barium     . 

.      Ba 

137-37 

Magnesium 

.     Mg 

24.32 

Bismuth  .      . 

.      Bi 

208.0 

Manganese 

.     Mn 

54-93 

Boron 

.      B 

II.O 

Mercury     . 

.     Hg 

200.6 

Bromine  . 

.      Br 

7Q.Q2 

Nickel  .     .     . 

.     Ni 

58.68 

Cadmium 

.     Cd 

II2.4 

Nitrogen     .      . 

.     N 

14.01 

Calcium  . 

.     Ca 

40.07 

Oxygen 

.     0 

16.0$ 

Carbon     . 

.     C 

12.  OO 

Phosphorus 

.     P 

31.04 

Chlorine  . 

.     Cl 

3546 

Potassium  . 

.     K 

39-1 

Chromium     . 

.     Cr 

52.0 

Silicon  .      ,     . 

.     Si 

28.3 

Cobalt      .      . 

.     Co 

58.97 

Silver    . 

.     Ag 

107.88 

Copper     .      . 

.     Cu 

63.57 

Sodium  . 

.     Na 

23.0 

Fluorine  . 

.     F 

I9.O 

Sulfur  .      .     . 

.     S 

32.06 

Gold  .      .      . 

v>     Au 

107.  2 

Tin.     .     .     . 

Sn 

118.7 

Hydrogen 

.     H 

V  /  * 

1.008 

Zinc      .     4 

.     Zn 

•-.****! 

65.37 

TENSION    OF   AQUEOUS   VAPOR   AT   VARIOUS    TEMPERATURES, 
EXPRESSED  IN  MILLIMETERS  OF  MERCURY 


TEMPERATURE 

15°  •  •  • 

16°  .  .  . 

17°  .  .  . 

18°  .  .  . 

19°  .  .  . 

20°  . 


PRESSURE  TEMPEKATURE  PRESSURE 

.  12.78  ^21° 18.62 

.  13.62  ^2°  .......  19.79 

.  14.52  23°  .    .    .    .   .    .   .  21.02 

.  15.46    24° 22.32 

.  16.56    25° 23.69 

.  17.51   100° 760.00 


WEIGHT  IN  GRAMS  OF  1  LITER  OF  VARIOUS  GASES  MEASURED 
UNDER  STANDARD  CONDITIONS 


Acetylene      .     .     .     .  1.1621 

Air 1.2928 

Ammonia       ....  0.7708 

Carbon  dioxide  .     .     .  1.9768 

Carbon  monoxide     .     .  1.2504 

Chlorine 3.1674 

Hydrogen      ....  0.08987 

Hydrogen  chloride    .     .  1.6398 


Hydrogen  sulfide  .     .     .  1.5392 

Methane 0.7168 

Nitric  oxide    ....  1.3402 

Nitrogen 1.2507 

Nitrous  oxide.      .      .      .  1.9777 

Oxygen 1.4290 

Sulfur  dioxide .     .     .     .  2.9266 


TABLE  OF  SOLUBILITY  OF  VARIOUS  SOLIDS 


WEIGHT  DISSOLVED  BY  100  cc.  OF  WATER  AT 

^tTR^TA  Wf"H 

« 

o° 

20° 

100° 

Calcium  chloride      •     . 

CaCl2 

59-5  g- 

74-5  g. 

JSQ-og- 

Sodium  chloride  .     .     . 

NaCl 

35-70  g. 

36-°  g- 

39.80  g. 

Potassium  nitrate      «     , 

KNO3 

1  3-30  g. 

3i-6  g- 

246.0  g. 

Copper  sulfate      .     .     . 

CuSO4 

14-3°  g- 

21.7  g. 

75-4g- 

Calcium  sulfate    .     .     . 

CaS04 

0-7  59  g- 

0.203  g. 

0.162  g. 

Calcium  hydroxide   .     . 

Ca(OH)2 

0.185  g- 

o.i65g. 

0.077  g- 

DISPLACEMENT  (ELECTROCHEMICAL)  SERIES 


1.  Potassium 

2.  Sodium 

3.  Lithium 

4.  Calcium 

5.  Magnesium 

6.  Aluminium 


7.  Manganese 

8.  Zinc 

9.  Chromium 

10.  Iron 

11.  Cobalt 

12.  Nickel 


13-  Tin 

14.  Lead 

15.  Hydrogen 

1 6.  Copper 

17.  Arsenic 

1 8.  Bismuth 


19.  Antimony 

20.  Mercury 

21.  Silver 

22.  Platinum 

23.  Gold 


[176] 


YB  36034 


568038 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


